Caveman Chemistry

Book 2

Experiments

15-28

Chapter 15. al-Razi (Stoichiometry)

Take one mann of white Al-Qilī and an equal quantity of Lime and pour over it [the mixture] 7 times its amount of water, and boil it until it is reduced to one half. Purify it [by filtration and decantation] 10 times. Then place it in thin

evaporating cups [kīzān], and hang it then in [heated] beakers [jāmāt]. Return what separates out [to the cup], and raise it [the cup] gradually, and protect

from dust whatever drops from the cups into the beakers, and coagulate it into salt.

— al-Rāzī, The Book of Secrets, ca. 920 AD [1]

15.1. Imagine that you are so pleased with your mead-making experience that you resolve

to brew up your entire 100 gallon/year entitlement at once. You'll have an entire year's worth of birthday presents, Christmas presents, anniversary and wedding gifts that you won't have to fret over. Who wouldn't enjoy a refreshing bottle of Samson's Special Reserve? You fit out two 55 gallon drums with fermentation locks to let the gas out without letting air in. You buy 100 pounds of honey, 100 pounds of lemons and 100 pounds of tea at wholesale prices and get a local bar to save 100 empty beer bottles for you. Not only don't they charge you, they even wash them in their industrial-size dishwasher. You realize that you can save on yeast since one yeast makes two and two make four and four make eight and so on, so you only need one packet of yeast. You're prepared for the yeast to take a little longer than usual to get going, but you're not expecting months to go by with precious little in the way of gas. Eventually you get tired of waiting and prepare to bottle the stuff anyway, but there's a problem.

Belatedly you realize that you don't have anywhere near the right number of bottles; you implicitly and erroneously assumed that each bottle would hold a gallon when in fact they are 12-ounce bottles. Hurriedly you use your old friend, UFA from Chapter 3 to calculate the number of bottles needed:

Before bottling, you give your Bee-jolet a little taste; it's terrible. More like industrial-strength lemon-tea furniture polish from Hell than anything you would even remotely consider drinking. That's when you realize that you made the same error with your ingredients that you made with the bottles. The poor yeasts never had a fighting chance!

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You spend the rest of the day flushing your experiment down the toilet and hoping that it doesn't set off a pollution alarm at the water-treatment plant.

You've probably already spotted your mistake, but let's walk through it anyway before you waste enough groceries that FoodWatch distributes "Wanted" posters with your smiling face on them. The honey calculation is easy:

Your original mead had less than half the honey called for by the recipe; no wonder it fizzled out.

The lemons are a little more difficult. You could use:

But your wholesaler sells lemons by the pound. A 10 pound bag of lemons contains an average of 40 lemons, so:

With almost twice the lemons called for by the recipe, it's no wonder your mead had more pucker than a goldfish in a pickle jar.

The tea holds complications of its own. You used tea in bags for your smaller batch, but your wholesaler sells bulk tea by the pound. A box of 25 tea-bags weighs 50 grams, so:

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Your original mead had 100 times the required amount of tea! It should have pleased any tea-totaler, since it was almost totally tea.

At this point in the narrative structure I'll bet you're wondering what your mead recipe can possibly have in common with lime and al-qilī. The Author wanted an example that would make absolutely clear the folly of confusing weight with number. You learned to balance metathesis reactions in Chapter 7 and redox reactions in Chapter 11. Balanced reactions give you the relative number of moles of each reactant and product but atoms and molecules are too small to count out individually. As in our mead example it is often more convenient to weigh things out than it is to count them out. The question of how much of one thing in a reaction goes with a given amount of another is called a stoichiometric question.

Notes

[1] Reference [26], p. 391.

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15.2. The chapter began with Muhammad bin Zakarīyā al-Rāzī's favorite recipe for making sodium hydroxide (caustic soda, lye) from sodium carbonate (al-Qilī, alkali, soda) and calcium hydroxide (slaked lime). This classic metathesis reaction ought to be familiar to anyone who managed to stay awake during Chapter 7. This recipe, one part soda with one part lime, brings up one of the most fundamental problems in all of chemistry; what, exactly, is a part?

In Chapter 7 I told you that the little numbers in front of each substance in a balanced chemical equation are called stoichiometric coefficients and that their unit is the mole. A mole is just a number like a dozen or a hundred or a million, but much bigger. Just as the weight of a dozen lemons would be expected to be different from the weight of a dozen tea-bags, the weight of a mole depends on what it is you're talking about.

The truth of the matter is that if you have a grip on Unit Factor Analysis , three new unit factors will allow you to solve any stoichiometric problem. And these three unit factors share a common unit, (grams/mole). The first of these is the atomic weight, the weight of one mole of atoms; the second is the formula weight, the weight of a mole of formula; similarly the molecular weight is the weight of a mole of molecules. Formula and molecular weights come from adding up the atomic weights for the elements in an empirical or molecular formula, respectively, and atomic weights may be found either in a periodic table or in Appendix E.

So let's jump right in with a problem from al-Rāzī:

Q: How many grams of soda can react with 100 grams of slaked lime?Q: How many grams of washing soda can react with 100 grams of quicklime?Q: How many grams of caustic soda can be produced from the reaction of 100 grams of quicklime with excess washing soda?Q: How many grams of anhydrous soda can be produced from the calcination of 100 grams of washing soda?Q: How many grams of mullite and silica can be produced from the calcination of 100 grams of dry kaolinite?

Q: How many grams of soda can react with 100 grams of slaked lime?

A: To answer any stoichiometry problem we must begin with a balanced equation:

From the balanced equation we get a mole ratio, in this case, for every mole of soda (sodium carbonate), one mole of slaked lime (calcium hydroxide) is consumed; one mole

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of limestone (calcium carbonate) and two moles of caustic soda (sodium hydroxide) are produced. Note that relative number of moles is given by the stoichiometric coefficient, the little number in from of each reactant or product in a balanced chemical equation. When there is no number in front of a reactant or product, it is assumed to be one. The unit factor required for this particular problem is (1 mole soda/1 mole slaked lime) or (1 mole slaked lime/1 mole soda), depending on the units we need in the top and bottom of the factor.

For the formula weights, we just add up the atomic weights; for soda, we get (2 +12 + 3ײ6ױ3 ) = (106 g/mole); for slaked lime we get (40 + 2 ױ2 + 6ױ ) = (74 g/mole).

That is, each part of slaked lime reacts with 1.43 parts of soda. So 1 part lime to 1 part soda may be almost 50% off, but it's not half bad for a recipe from the tenth century.

Remember that a string of unit factors can be continued from one line to the next and that a second "=" sign begins a second equation which is equal to the first. You could easily have worked this problem in Chapter 3 if you had been given the unit factors in the statement of the problem. The only new things in this chapter are the formula weights and the mole ratios.

Q: How many grams of washing soda can react with 100 grams of quicklime?

A: So far, I have used the terms soda ash and washing soda interchangeably. When either one dissolves in water it immediately dissociates into sodium ions and carbonate ions. Soda dissolved in water has no "memory" of where it came from, so chemically speaking, anhydrous (without water) soda and washing soda can be used interchangeably. Nevertheless, they are not exactly the same substances; in particular, their formula weights differ and this is relevant to the problem of how much to weigh out to achieve a balanced reaction. Similarly, as soon as quicklime hits the water, it slakes and the resulting slaked lime is identical to the slaked lime of the previous problem. But the weight of lime we should use depends very much on whether we are using quicklime or slaked lime. When weighing out chemicals, it is vital to know whether the material is hydrous or anhydrous because you must account for the weight of any water present.

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The formula weights of washing soda and quicklime are 286 g/mole and 56 g/mole, respectively.

About 5 parts of washing soda are required to react with each part of quicklime.

Q: How many grams of caustic soda can be produced from the reaction of 100 grams of quicklime with excess washing soda?

A: There are two wrinkles in this one; the mole ratio is no longer 1:1. The word excess tells us that there is "more than enough soda". We assume that all of the quicklime is consumed and that there is leftover soda when the reaction is done. The reactant that runs out first, in this case quicklime, is known as the limiting reagent, the one which is not in excess. For now, the word excess tells you that you only need to deal with the other reactant.

You may be disturbed by the notion that you can make 146 g of caustic soda from only 100 g of quicklime. Is the Law of Conservation of Mass violated here? Not at all! If you calculate the number of grams of soda reacting with that 100 grams of quicklime, you will see that the weight of the soda plus the weight of the quicklime is exactly the same as the weight of the caustic soda plus the weight of the chalk plus the weight of the water produced.

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Q: How many grams of anhydrous soda can be produced from the calcination of 100 grams of washing soda?

A: This would be a very easy experiment to perform. Just put a weighed amount of washing soda into a hot oven (above 100?C) for a few hours and weigh it again when it is done. Of course, you would substitute your actual weight of washing soda for the "100 grams" in the stated problem.

The formula weights of washing soda and anhydrous soda are 286 g/mole and 106 g/mole, respectively.

Your calculated weight of product, in this case 37 g of anhydrous soda, is called the theoretical yield. It can be compared to your actual weight, the experimental yield. When the theoretical and experimental yields are close to one another, you have a happy, comfortable feeling that you understand what's going on and that you are getting the most from your reaction. If they are different from one another, you have an uncomfortable feeling that you don't understand what's going on, and that you are not getting everything from your reaction that you might. Think about some situations that might result in the theoretical yield being higher or lower than the experimental yield. Then try it for yourself.

Q: How many grams of mullite and silica can be produced from the calcination of 100 grams of dry kaolinite?

A: Now, you were supposed to have weighed your pot before and after firing, and so at this point you can see for yourself whether Athanor was feeding you a line.

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There are several conditions assumed by the question itself. The calculation assumes that we started with pure, dry kaolinite, a condition that may have been only approximated in your clay body. It assumes that Equation 5-1 is the only reaction taking place and that all of the kaolinite was converted, that is, no un-reacted kaolinite remains in your fired pot. And it assumes that nothing stuck to or flaked off of your pot. Substituting the weight of your un-fired pot for the "100 grams" of the problem, calculate a theoretical yield and compare it to the experimental yield. Of course, you have to calculate the theoretical weights of mullite and silica separately and then add them to get the theoretical weight of your pot. If your theoretical and experimental yields are close to one another, it doesn't necessarily mean that everything Athanor told you was true, it just means that your evidence supports his assertions. Scientific theories are never proven once and for all; but the more evidence supports a theory, the more confident we become that it is essentially correct.

Stoichiometry in a NutshellChill out, whatcha yellin' for? Lay back, you've done it all before.Put down the grams of stuff you want to find, and an equal sign.Then put what you're startin' from. You don't have to play so dumb.I mean the problematic units, hey, like UFA.

The formula weight of the thing that you hatewill cancel the gram like an unwanted spam.

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A mole ratio takes you where you gotta go, you see. Tell me.

Why d'you have to make stoichiometry complicated?Don't you know itdoesn't have to be the kind of thing that leaves you all frustrated?Life's like this youcancel moles that you hate, one more formula weight,gives you grams of the answer, on toJust a little number-crunchin' work and you know you made it.Yea, yea, yea.

Figure 15-1. The NFPA Diamond

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Material Safety

If you're like most people you find the information in an MSDS rather daunting. It's full of all kinds of technical information suited to occupational safety. Wouldn't it be great if there were some kind of simple, easily understood safety rating system so that you would know whether a chemical was really, really dangerous, or just dangerous in the way that rocks and sticks and dirt are dangerous? Well, there is such a system. You've probably seen symbols like those in Figure 15-1 which provide information on hazardous materials, but most people don't have a clue what they mean.

The National Fire Protection Agency[1] developed the NFPA diamond to give fire fighters information on hazardous chemicals at a glance. The number in the left-most (blue) quadrant is the health rating, that in the upper-most quadrant (red) is the flammability rating, that in the right-most (yellow) quadrant is the reactivity rating, and the number in the lowest (white) quadrant gives special hazards, for example, OX for oxidizer or W for exceptional reactivity with water. The numerical ratings run from 0 (not very dangerous) to 4 (very dangerous). Most MSDS's give the NFPA ratings, either as a diamond graphic or as text (NFPA H 1, F 0, R 0, OX). Let's look at the NFPA diamonds (Figure 15-1) for four materials which will play a large role throughout the remaining chapters.

We've already been introduced to ethanol in Chapter 4 and will see more of it in Chapter 16. As a poison, ethanol is fair to middling and gets a "2" in the health quadrant. In the flammability quadrant ethanol gets a "3;" it's quite flammable, though not as flammable as something like propane. In the reactivity quadrant ethanol gets a "0;" it's not particularly reactive with common materials and is not explosive. The special quadrant is blank, as ethanol isn't an oxidant and it isn't particularly reactive with water. Ethanol's NFPA symbol tells a firefighter that flammability is its most important property in an emergency situation.

We met caustic soda, sodium hydroxide, in Chapter 14 as an excellent alkali for making paper. Sold in the grocery store as "lye," it's a common ingredient in household drain openers. Caustic soda will emerge as an important player in Chapter 19. In the present chapter we've explored the stoichiometry for making caustic soda from soda and lime. Caustic soda is fairly toxic by ingestion, but the fact that it eats skin bumps its health rating up to a "3." Caustic soda will not burn, so it gets a "0" for flammability. It is reactive with some other materials, notably aluminum, and so it gets a "1" for reactivity. While it gets hot on contact with water, this is not enough to get it a notice in

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the special quadrant. Thus caustic soda's NFPA symbol tells the firefighter that exposure to the material should be avoided, but that it is neither flammable nor explosive.

Ordinary laundry bleach is a solution of sodium hypochlorite in water. It is roughly as poisonous as caustic soda and gets a "3" for health. It's not flammable so it gets a "0" for flammability. It's not particularly reactive so it gets a "0" for reactivity. It is an oxidizer, meaning that it can substitute for oxygen in a fire, a fact to which fire fighters are alerted in the special quadrant.

We'll meet nitroglycerin in Chapter 27, though you probably know it already as a powerful explosive. While it's moderately toxic and flammable, its most important property from a firefighting point of view is that it will blow you to kingdom come. That gets it the highest rating, a "4," in the reactivity quadrant.

The diamond was designed for fire fighters and the NFPA doesn't encourage its use for other purposes. The diamond, for example, doesn't alert people to chronic toxic effects or carcinogenicity. Thus it's less relevant to hazards associated with long-term occupational exposure than it is to acute exposure under emergency situations. At least two other rating systems have emerged to deal with occupational and laboratory situations: the HMIS rating of the National Paint and Coatings Association[2] and the Saf-T-Data rating of the J. T. Baker chemical company.[3] Like the NFPA diamond, both systems use a numerical rating from 0 (not hazardous) to 4 (very hazardous) in four or five categories akin to those of the diamond. Unlike the diamond, however, these systems factor the effects of chronic exposure into the health rating and so the numerical ratings may differ somewhat from one system to another. Either the HMIS or Saf-T-Data systems would appear to be preferable to the NFPA diamond for laboratory use, but neither has been adopted as widely as the diamond. Since the diamond is nearly universal and since chemical exposure for cavemen is likely to be acute rather than chronic, we'll concern ourselves primarily with the diamond as a shorthand for chemical hazards. Still, it's only a shorthand and not a substitute for the full MSDS.

You aren't going to get away without a material safety assignment, even though there are no materials involved in this project. If you don't have them already, search for MSDS's using the keyword "NFPA." Draw the NFPA diamonds for slaked lime (CAS 1305-62-0), propane (CAS 74-98-6), and silica (CAS 14808-60-7) in the material safety section of your write-up for this project.

Research and Development

So there you are, studying for a test, and you wonder what will be on it. Study the meanings of all of the words that are important enough to be

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included in the index or glossary. Know the Research and Development items from Chapter 9 and

Chapter 11. Practice calculating formula and molecular weights from atomic

weights. Most people don't expect you to memorize atomic weights nowadays.

But you had better be able to find them on a periodic table. Just remember that the atomic weight of carbon is 12. The atomic weights of the other elements will appear in the same position that 12 does for carbon.

All of the problems here have used 100 grams, so you are getting a percentage, but you should expect to be able to solve problems using different amounts.

Understand the meaning of the NFPA diamond.

When you can't remember the last time you missed a stoichiometry problem, you've probably worked enough examples.

Notes

[1] For more information on the NFPA diamond, see Reference [35].[2] For more information on the HMIS system, see Reference [37].[3] For more information on the Saf-T-Data system, see Reference [31].

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15.3. Q: How many grams of quicklime can be produced by the complete calcination of 100 grams of limestone?Q: How many grams of water are needed to completely slake 100 grams of quicklime?Q: How many grams of slaked lime can be produced from the reaction of 100 grams of quicklime with excess water?Q: When cellulose is heated in a reducing environment it becomes charcoal as shown in Equation 1-1. How many grams of charcoal can be produced by the complete conversion of 100 grams of cellulose?Q: The partial calcination of gypsum produces plaster according to Equation 10-1 . How many grams of plaster can be produced by the complete conversion of 100 grams of gypsum?Q: How many grams of copper can be produced by the complete reduction of 10 grams of copper carbonate?

Q: How many grams of quicklime can be produced by the complete calcination of 100 grams of limestone?

A: 56 grams of quicklime.

Once again, you can compare the theoretical and experimental yields from your lime-burning experience. Be alert for the possibility that your limestone was impure or your calcination incomplete.

Q: How many grams of water are needed to completely slake 100 grams of quicklime?

A: 32 grams of water.

Q: How many grams of slaked lime can be produced from the reaction of 100 grams of quicklime with excess water?

A: 132 grams of slaked lime.

Q: When cellulose is heated in a reducing environment it becomes charcoal as shown in Equation 1-1. How many grams of charcoal can be produced by the complete conversion of 100 grams of cellulose?

A: 40 grams of charcoal.

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Q: The partial calcination of gypsum produces plaster according to Equation 10-1. How many grams of plaster can be produced by the complete conversion of 100 grams of gypsum?

A: 84 grams of plaster.

Q: How many grams of copper can be produced by the complete reduction of 10 grams of copper carbonate?

A: 5.1 grams of copper. Of course, you can go further and ask how many grams of tin will be produced by the complete reduction of 2 grams of tin oxide; the theoretical yield of bronze can then be had by adding those of copper and tin. The usual conditions apply; perhaps what you thought was pure copper carbonate was actually pure or impure malachite, a hydrate of copper carbonate; perhaps more than one reaction was going on in the kiln; perhaps you were unable to collect all of the metal produced; perhaps charcoal was the limiting reagent and consequently you have some un-reduced ore and less metal than you would expect. When the theoretical and experimental yields agree, however, you get that happy feeling that you understand what is going on and that all is right with the world.

Quality Assurance

Tape your stoichiometry quizzes into your notebook and calculate the theoretical yield for one of your previous projects: pottery, lime, gypsum, or metal. Compare the theoretical yield to your experimental yield. Speculate on the kinds of things that could lead to any discrepancies.

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Chapter 16. Adelard (Alcohol)

Since I possess many wonderful books written on these matters I became anxious to produce a commentary, not that I may appear to be encroaching upon the sacred books and [therefore] despite much labor accomplishing nothing, but that, avoiding that mortal heresy, I will disclose to those who wish to understand these things what the actual processes are that are used in painting and other kinds of work. I call the title of this compilation Mappae Clavicula, so that everyone who lays hands on it and often tries it out will think that a kind of key is contained in it. For just as access to [the contents of] locked houses is impossible without a key, though it is easy for those who are inside, so also, without this commentary, all that appears in the sacred writings will give the reader a feeling of exclusion and darkness. I swear further by the great God who has disclosed these things, to hand this book down to no one except to my son, when he has first judged his character and decided whether he can have a pious and just feeling about these things and can keep them secure.

…

212. From a mixture of pure and very strong xjnf with 3 qbsut of tbmu, cooked in the vessels used for this business, there comes a liquid, which when set on fire and while still flaming leaves the material [underneath] unburnt.— Mappae Clavicula, ca. 1130 AD [1]

16.1. This flaming liquid is one of the most remarkable substances ever discovered and so

you will understand that I have endeavoured to hide it among my secrets. It is nothing less than the vegetable mercury, which rises from the earth to heaven and descends again to the earth. It is the strong power of all powers for it overcomes everything fine and penetrates everything solid. In Latin it is called Aqua Vitae, in Gaelic Uisce Beatha, both names meaning "water of life." The English alcohol is derived from the Arabic al-Kuhl. My child, are ye prepared to preserve this immortal fire, to protect it from the ignorant rabble, to keep in its heat and withstand it?

Hey, I was sitting there.

You are probably wondering what is going on. I will tell you. I was minding my own business, preparing to tell you about alcohol, it being my turn to speak. I stepped out to answer the memetic equivalent of the call of Nature and returned to find Lucifer in the driver's seat preparing to make off with my chapter. Now I admit that my subject touches on heat and temperature and winds up with a fondue; still, if Lucifer were a chicken she would not belong in every pot. But I am getting ahead of myself.

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Of course, alcoholic beverages have been around since God was a child. The earliest written records speak of mead, wine and beer as part of everyday life. But the isolation of alcohol from these sources is of much more recent invention. You see, while any old devil can heat the bejeezus out of something, it is quite another matter to catch this bejeezus and put it in a bottle. Catching a spirit requires a kinder, gentler kind of fire, and once the bejeezus has been coaxed from its lair it must be herded like a cat, not summoned like a dog.

People will argue about who was the first to succeed in distilling alcohol. I can only tell you that I myself inhabited a painter of the twelfth century when I discovered it. Now, a painter had to be a jack of all trades in those days. One day you might be painting with oils on canvas, another dyeing yarn for a tapestry, another gilding letters in a book, and yet another preparing plaster from gypsum. And because there was quite a bit of variation in the quality and purity of the various metals, dyes, pigments, binders and solvents an artist really had to be something of a chemist if he was not to be taken to the cleaners.

Even more-so, an artist had to be on the lookout for new combinations of colors and textures to provide for the endless struggle of queens and kings to out-dress call-girls and clowns. In the dog-eat-dog world of high fashion, a successful painter had to jealously guard his recipes from theft by his competitors. Now in those days there were many solvents for painting and dyeing; water for some, oil for others, wine, vinegar or urine for still others, and so on. It was difficult enough finding a solvent for one pigment, let alone trying to mix a water-soluble pigment with an oil-soluble one. What was needed was a universal solvent, one that would dissolve just about anything.

Metallurgy had such a solvent in mercury. You see, quicksilver will dissolve most any metal, particularly the valuable ones like silver and gold. In fact, one method for recovering gold from it ore is to mix it with mercury, which dissolves the elemental gold but not the oxides, sulfides, and carbonates of the other metals. This mercury, laden with dissolved gold, is then placed in a furnace in which the bejeezus is heated from it, that is, the mercury is boiled away leaving gold behind. Now it had been noted in antiquity that when mercury is boiled it disappears from the pot, but, amazingly, beads of mercury collect on any cool surfaces which may happen to be nearby. Because mercury does not grow on trees, it was advantageous to arrange matters so that there was always a cool surface nearby so that the mercury might be recovered. Such an arrangement came to be known as a still.

So the idea was instilled in me that mercury, the universal metallic solvent, was a kind of spirit which could be driven out of a metal by heating and collected by cooling. The same was true of another very good solvent, water. It became my habit, then, in the odd moment when nothing else was cooking, to distill whatever leftovers happened to be lying about to see what might come of them. Many people had the same notion and wine, in particular, was a popular choice for distillation. The resulting "spirits" provided an enhanced kick, but were not qualitatively better solvents than wine.

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One day I happened to throw some salt into the pot for no particular reason. The resulting distillate was stronger than any spirit I had ever seen. Now, taste is an important tool for chemical analysis when you have nothing else and so I had a little shot of my new spirit, purely in the spirit of scientific inquiry you understand. It literally knocked me off my feet, with the result that I lost my balance and dropped my beaker near the furnace. Whoosh! The whole place erupted in flames and I thought sure I had fallen in the clutches of the fellow with the horns. Fortunately the fire burned out as quickly as it started and I was left a little singed but a lot wiser.

I suppose that the lesson most people would have taken from this experience is never to mix salt with wine. But I was not most people. No, the lesson I learned was to sit before sipping solvent. Eventually my attention returned to matters artistic and I found that this new spirit was an almost universal solvent for pigments, dyes and medicines. When I wrote down the recipe for this phenomenal potion, I did so in code, lest I lose my competitive advantage. You are probably wondering about the secret of this code. I would tell you but then it would no longer be a secret.

As it turned out, my code was soon broken. You see, it is in the interest of the mortal to keep secrets, but it is not in the interest of secrets to be kept. Memes which do not get out much are apt to go extinct and so having wriggled my way out of my medieval painter, it is I, the un-kept secret, who remain to write this chapter. So I am now prepared to let the alcoholic cat out of the bag, so to speak.

Notes

[1] Reference [25].

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16.2. First of all, I must tell you that when a less volatile solute is dissolved in a more volatile solvent, the boiling point of that solvent goes up. You probably imagine that volatile means flammable or unstable or explosive. If so you imagine incorrectly. No, in chemistry to say that one compound is more volatile than another simply means that it is easier to evaporate, that its boiling point is lower. So, for example, the boiling point of salt-water is higher than that of water. The boiling point of mead is higher than that of alcohol and lower than that of water.

Imagine if you will, a tea-kettle, called by chemists a pot on the stove filled with mead. The heat is on and the mead comes to a boil. If it were pure alcohol it would boil at 78?C; if it were pure water it would boil at 100?C. But because it is a solution of alcohol in water, it boils somewhere in between, let us say, at 90?C. Now if we were to collect a bit of the steam coming out of the spout, what do you suppose we would find? The alcohol "wants" to boil at 78?C so at 90?C it is really ready to fly the coop, so to speak. By contrast, the water does not really "want" to boil until 100?C so at 90?C it is in no hurry to bolt. Put another way, because alcohol is the more volatile of the two, the steam will be richer in alcohol than the original mead.

Imagine further that you place a cold dinner plate over the spout of the tea-kettle. The steam will condense into droplets of liquid on the cold surface. This liquid will have the same concentration of alcohol as the steam, that is to say, it is richer in alcohol than the original mead. We could collect this liquid and place it into another tea-kettle, albeit a tiny one. Being richer in alcohol, it would come to a boil at a lower temperature than mead, perhaps 85?C. You are probably thinking that we could then collect steam from this new kettle, condense it on a second cold plate, and dribble this even richer liquid into a third kettle. If so, your idea would be gathering steam.

But it is not actually necessary to use such a cumbersome arrangement. If, instead of a dinner plate, we mounted a tube, called by chemists a column on the spout of the tea-kettle, the steam would condense on the inside. As the bottom of the tube gets hot from the rising steam, the drops of condensate will boil again, as it would have in a second kettle. The new steam rises a little further up the tube until it finds a place cool enough to condense and the whole process starts all over again. A simple tube, hot at the bottom and cooler at the top, is all that is needed to take the place of a whole brigade of tea-kettles and the longer the tube is, the richer in alcohol the distillate will be.

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Figure 16-1. Distillation as a Process

Imagine further that we attach another tube, a condenser to the top of the column. The only real difference between the column and the condenser is that the column is vertical, so that condensing liquid dribbles back down toward the pot while the condenser has a downward slant so that condensing liquid can exit the still. The condenser might be cooled by air, but it is more efficient if it is cooled by water. The place where column and condenser meet is called the head. Imagine now a thermometer at the head. What would it read during a distillation? At the beginning there is boiling mead in the pot but the head is still at room temperature. The steam begins to rise, heating the bottom of the column, but the head remains at room temperature. Eventually, though, hot steam reaches the head where it simultaneously passes into the condenser. If we were distilling pure alcohol, the thermometer would read 78?C and the temperature would remain constant throughout the distillation. If we were distilling pure water, the temperature would remain at 100?C throughout. But with a solution, the head temperature can be anywhere in this range, lower when the distillate is more alcoholic, higher when it is less so. Unlike pure substances, which boil at a single, fixed temperature, solutions boil over a range of temperatures and this range can be used to monitor the concentration of the solution.

Actually, there may come a point where the boiling point of the distillate is lower than that of the alcohol. The steam coming off of this distillate will be no richer in alcohol than the liquid and so further distillation is ineffective. For ethanol-water this azeotrope boils at 78.17?C, 0.23?C lower than pure alcohol, and consists of 95% ethanol and 5% water. "Denatured" alcohol is simply 95% ethanol with little bit of poison added to discourage people from drinking it.

Figure 16-1 shows the distillation process in schematic form. The first reactor looks like a furnace, but one which contains liquid instead of solid. The second reactor is the condenser. Together, they make up a still. No chemical reaction takes place in this process. The process physically separates two liquids that differ in boiling point. The usual conventions are followed; reactants enter from the left of the figure, waste products exit the top and bottom, and the main product exits to the right.

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Adelard's motivation for distilling alcohol was to use it as a solvent and so I should probably tell you a bit about solubility. There are essentially three kinds of substance; ionic substances, or salts consist of positive and negative ions and were discussed in Chapter 7; polar substances consist of molecules in which the negative charge is bunched up at one end and the positive charge at the other; non-polar substances consist of molecules in which the charge is evenly distributed throughout the molecule. In general, substances in which the charges are separated, salts and polar substances, are mutually soluble and non-polar substances are mutually soluble, but salts and polar substances are not soluble in non-polar substances and vice versa. "Oil and water do not mix," or, to put it another way, "like dissolves like."

Figure 16-2. Two Water Molecules

Polar molecules, as I have said, have negative charge bunched up at one end and positive charge at the other. For this to happen, one or more atoms must, like the devil, take more than its due. Of electrons, that is. While carbon and hydrogen are rather mild-mannered in this regard, oxygen and nitrogen are greedy. In the water molecule, for example, the oxygen atom hogs electrons from the hydrogen atoms and so the oxygen end is negative and the hydrogen end is positive. When two water molecules bump into one another, as shown in Figure 16-2, the negative oxygen end of one molecule is attracted to the positive hydrogen end of the other. Opposite charges attract, you see. This mutual attraction, or hydrogen bond, is what holds two water molecules together. A large collection of water molecules resembles a dog convention, with the little beggars sniffing at each other's tails.

Ethane, C2H6, is a typical non-polar molecule. Because carbon is really no greedier than hydrogen, as far as electrons are concerned, there is no positive or negative end to the molecule and hence no strong interaction to bind one ethane molecule to another. In contrast to water, ethane molecules are content to roam about passing one another like head-less tail-less dogs in the night.

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What happens when a polar liquid like water and a non-polar one like ethane find themselves in the same container? The situation is much like a room filled half with normal dogs and half with head-less, tail-less dogs. The normal dogs will congregate in one part of the room because of their mutual attraction. The headless, tail-less dogs will be pushed to the other side of the room, not so much because they are attracted to one another, but because the normal dogs are less attracted to them than to those of their own kind. So to come back to chemistry, the polar water molecules will collect in one part of the container and the non-polar ethane molecules will be left in the other. To put it another way, ethane is not very soluble in water and vice versa.

Of course, the question of polarity is not black-and-white. There are some molecules that are very polar, some that are a bit polar, and some that are not very polar at all. For example ethanol, C2H5OH, is mid-way between ethane and water in its polarity. One end of the molecule is non-polar like ethane while the other is polar like water. It is as if a normal dog had been sewn onto a head-less tail-less dog. Such a dog, released onto a pack of dogs, would be equally comfortable with either kind. In Figure 16-3 a water molecule is attracted to the oxygen-hydrogen end of the ethanol molecule and vice versa. This mutual attraction is what is responsible for the ability of ethanol to dissolve in water.

Figure 16-3. Ethanol and Water

You may be wondering why ethanol is such a good solvent for so many dyes and medicines. If you think about it, living things, whether yeasts or bacteria, trees or polliwogs, are made up of tiny bags of water called cells. Part of the living thing is polar, the part dissolved in the water. But much of the organism had better be non-polar, the cell walls, skin and bones, or else the whole creature would dissolve in its own juices and be reduced to a blob of goo. Organisms need to move things around and these things must be polar, but once moved to where they belong they need to stay put. So living things are constantly manipulating their components, adding OH's here and there to get them into solution and then chopping them off or covering them up to bring them out again. The beauty of ethanol, or ethyl alcohol, as a solvent is that it can dissolve many organic compounds whether or not they are soluble in water.

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I should mention that ethanol is only one member of a class of compounds, the alcohols.Methanol, CH3OH, is also known as wood alcohol since it came originally from the distillation of wood. Isopropanol, C3H7OH, is familiar as rubbing alcohol. Ethylene glycol, C2H4(OH)2, which looks like ethanol with an OH group on each carbon atom, is used as automobile anti-freeze. Glycerol, C3H5(OH)3, like isopropanol with an OH on each carbon, will be discussed at length in Chapter 19, when we discuss soap. Each of these compounds consists of a carbon-hydrogen chain, the non-polar bit, with one or more polar OH groups hanging off of it.

Figure 16-4. Acetic Acid and Water

When ethanol is oxidized by bacteria (Equation 16-1(a)) it becomes acetic acid, which makes up about 5% of vinegar. Acetic acid, CH3COOH, is one member of the class of compounds, the organic acids. You will recall that oxygen is an electron hog and with two of them pulling electron density away from the poor hydrogen, it is left with a greater positive charge than it had in either ethanol or water. Since electrons are what hold the molecule together, this acidic hydrogen or proton has a tendency to abandon ship if it happens to bump into an alkali. This kind of metathesis reaction, illustrated in Equation 16-1(b), is characteristic of an acid-base reaction; the acid donates a proton to the base.

Equation 16-1. Reactions of Ethanol and Acetic Acid

Now, like sodium hydroxide, ethanol has an OH group and you are possibly wondering whether a metathesis reaction is possible between ethanol and acetic acid. If so, you are wondering in the right direction. Equation 16-1(a) shows the metathesis reaction producing ethyl acetate, which is a member of the class of compounds, the esters. Esters are the organic counterparts of inorganic salts, being produced as a condensation of an

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acid and an alcohol. If your mead contained bacteria and oxygen you no doubt produced some ethyl acetate by accident. Unlike alcohols and acids, however, esters have no OH groups and the negative oxygen atom is buried deep inside the molecule. Consequently ethyl acetate is not very polar and makes an excellent solvent for non-polar compounds. It is a popular choice for fingernail polish remover.

Figure 16-5. Ethyl Acetate

I have still to tell you what salt, or, as I like to call it, tbmu, has to do with anything. In those days our condensers were short, air-cooled jobbies, good enough for separating substances like gold from mercury, whose boiling points differ by hundreds of degrees. I did not suspect, and was not prepared to efficiently separate substances like alcohol from water, whose boiling points differ by only 22 degrees. As I have said, when a less volatile solute is dissolved in a more volatile solvent, the boiling point of that solvent goes up. Since salt dissolves in water but not in ethanol, it raises the boiling point of the water but not the ethanol, making their boiling points farther apart than they would normally be and consequently making it possible to separate them even with an inefficient air-cooled condenser. When you do not suspect that such a thing is possible, there is, of course, no reason to worry about condenser efficiency, but after my discovery it became clear that most of the alcoholic genie was going out the window rather than into the bottle. And while the painter himself tried to keep me bottled up, I was too volatile for him to keep a lid on me; I spread my bejeezical wings and flew the coop shortly after he gave up the ghost.

Material Safety

Locate an MSDS for ethanol (CAS 64-17-5). Summarize the hazardous properties of this material in your notebook, including the identity of the company which produced the MSDS and the potential health effects for eye contact, skin contact, inhalation, and ingestion. Also include the LD50 (oral,

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rat) and the NFPA diamond for ethanol.[1]

While it is possible to drink enough beer, wine or mead to make yourself seriously ill, it is even easier with distilled spirits. An eight-ounce glass of ethanol is the equivalent of more than a half-gallon of wine or a gallon of beer. Remember, it is the dose that makes the poison. You should also be aware that ethanol dissolves things that are not soluble in more dilute alcoholic solutions. Consequently, poisonous compounds from bottles or stills may dissolve in ethanol that would not have dissolved in mead. Finally, you should be aware that ethanol is not the only volatile compound present in fermented beverages so when you distill alcohol you may concentrate other potentially toxic compounds, ethyl acetate, for example.

You should wear safety glasses while working on this project. Your distilled spirit may be carefully burned or saved for use in a spirit lamp, as shown in Figure 16-8. The contents of the pot may be flushed down the drain.

Research and Development

You are probably wondering what will be on the quiz. You should know the meanings of all of the words important enough to

be included in the index or glossary. You should have studied the Research and Development items from

Chapter 11 and Chapter 12. You should know the reactions of Equation 16-1. You should know the name and purpose of each part of a still. You should know why ethanol is such a good solvent for both polar

and non-polar compounds. You should know the hazardous properties of ethanol.

You should know why secrets are so hard to keep.

Notes

[1] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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16.3. The 18th amendment to the United States Constitution prohibited "the manufacture, sale, or transportation of intoxicating liquors within, the importation thereof into, or the exportation thereof from the United States." The law was so popular that it was repealed thirteen years later. The current law in the United States allows each adult to brew up to 100 gallons of beer or wine for her own personal use and a home-brewing industry has grown up to service hobbyists. It remains absolutely illegal, however, to distill spirits without a government permit. Boot-leggers are enthusiastically hunted down by law enforcement officers who, like the tenth-century painter, want to keep a lid on the spread of spirits.

Still, Part 19.901 of the Code of Federal Regulations "implements 26 U.S.C. 5181, which authorizes the establishment of distilled spirits plants solely for producing, processing and storing, and using or distributing distilled spirits to be used exclusively for fuel use." To establish such a plant requires a simple alcohol fuel producer's permit, form 5110.74, which may be ordered from the Bureau of Alcohol, Tobacco, and Firearms.[1] There is no fee for the permit and no tax is levied on alcohol used as fuel. Do not imagine that such a permit allows you to drink your spirits. If you are going to break the law anyway, such a permit is of no use to you. But if you are driven by curiosity to experience first-hand one of the most marvelous, literally spiritual substances in all of chemistry, this "liquid, which when set on fire and while still flaming leaves the material [underneath] unburnt," such an easily-acquired permit allows you to do so without getting into hot water with the Feds.

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Figure 16-6. The PVC Still

Of course, you will need a still and, there being no home-distillery market, you will have to either buy a still intended for laboratory use or build one yourself. Fortunately a still is not a particularly complicated piece of apparatus. I will tell you how to build one from an Erlenmeyer flask, a rubber stopper, and some common plumbing supplies. The complete still is shown in Figure 16-6(L) and the pot is detailed in Figure 16-6(R). Part (a) is a rubber stopper sized to fit the 500 mL Erlenmeyer flask, part (b). The stopper should have a hole drilled in it, 5/8 inch in diameter. Erlenmeyer flasks are not expensive and you can get one with a stopper from one of the suppliers listed in Appendix D. A flask weight will help prevent the still from tipping over. The one shown is constructed from four 3-inch sections of half-inch copper tubing, parts (c), and four elbows, parts (d), available from a plumbing supply. Fill the tubing with lead or steel shot before gluing it together with epoxy cement. The weight fits over the neck of the Erlenmeyer flask, which sits atop a hot-plate.

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Figure 16-7. Details of the Column and Condenser

The column, detailed in Figure 16-7(L), is constructed from two 9-inch sections of half-inch CPVC tubing, parts (e), a common meat thermometer, part (f), two half-inch elbows, parts (g), and a section of tubing, part (h), long enough to accommodate the thermometer; the one shown is 4 inches long. Be sure to use CPVC tubing—the kind intended for plumbing hot water. Drill a hole in one of the elbows to make a tight fit with the probe of the meat thermometer. Use cement specifically designed for CPVC to glue the column together.

Any convenient container will serve as the receiver, part (i). The one shown in the figure is a 20-ounce soft drink bottle. The condenser, part (j), consists of the bottom of a 2-liter soft drink bottle filled with crushed ice. Commercial condensers are air-cooled or water-cooled, but the one shown works admirably without increasing the cost or complexity of the still.

You need a source of heat strong enough to boil water but not so intense as to crack the Erlenmeyer flask. Laboratories are equipped with hot-plates designed for exactly this purpose. Alternatively, there are many varieties of consumer appliances such as hot-plates, coffee makers, and crock pots which could be made to serve. Avoid anything whose heating elements visibly glow, as it may cause the glass to crack. It is a good idea to keep a leather work glove on one hand during the distillation so that you may quickly remove the still from the hot-plate if the temperature rises too quickly.

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To operate the still, fill the Erlenmeyer flask with 400 mL of mead, wine, or beer. Add a couple of "boiling chips," small pieces of glass or silica, to prevent the still from boiling over. Slip the flask weight over the flask and insert the rubber stopper and column into the flask. Turn on the hot-plate, gradually turning up the power until the solution in the flask comes to a boil. Place the receiver into the condenser, fill the condenser with crushed ice, and slip the free end of the still into the receiver. The free end of the column should go all the way to the bottom of the receiver.

You are about to learn the truth of the old saw, "a watched pot never boils." You will watch the thermometer for some time with no change in its temperature. The temperature will begin to rise only when hot vapor makes it up the column to the head, the elbow where the thermometer lives. Do not get impatient—if you turn up the heat too quickly, the head temperature will be too high and your distillation will be spoiled. Eventually, the head temperature will rise and "steam" will exit the column into the receiver, where it will condense. Because the boiling point of ethanol is lower than that of water, your goal from this point on is to maintain the head temperature at the lowest possible value while still collecting liquid in the receiver. The boiling point of ethanol is 78?C (172?F); if you can keep the head temperature at or below 83?C (181?F) you will get reasonable separation of ethanol from water. If the head temperature gets a little too high, reduce the power to the hot-plate. If it is way too high, use your gloved hand to remove the still from the hot-plate and allow it too cool for a minute or so. After a while you will get into a "zone," in which the head temperature remains constant and liquid accumulates in the receiver. At some point you may wonder whether anything is coming out of the column; simply remove the receiver and cautiously feel the end of the column with your non-gloved hand. If it is hot, you are still collecting alcohol. Eventually you will find that you cannot continue collecting alcohol without increasing the head temperature. You will have collected most of your alcohol in the receiver and the pot will contain mostly boiling water and yeast carcasses. You may now turn off the hot-plate and allow the still to cool. Disassemble the cold still and pour the pot liquid, but not the boiling chips, down the drain. Rinse out the column and put it away. You may screw the cap on your receiver and store your alcohol for later use.

A crude estimate of the alcohol content can be made by burning your alcohol distillate. Weigh a Petri dish and record its empty weight, wbefore in your notebook. Tare the balance and add 10.0 g of your distillate, wdist, to the dish and set it alight. When the fire has gone out, weigh the Petri dish again and record this weight, wafter in your notebook. The percentage of alcohol of your distillate is approximately:[2]

[95 - 100(wafter-wbefore)/wdist]%

Double the percent alcohol and you have the proof. If your product is below about 40% alcohol (80 proof), you will be unable to light it at all.

Do not overlook this experience, my brothers and sisters. To merely read and parrot the words is unworthy of your talents. If you are to really live, you must see and feel and know for yourself.

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It rises from the earth to Heaven and descends again to the Earth and receives power from Above and from Below. Thus thou wilt have the glory of the Whole World. All obscurity shall be clear to thee. This is the strong power of all powers for it overcomes everything fine and penetrates everything solid.— The Emerald Tablet of Hermes Trismegistos

Scat!

Figure 16-8. The Spirit Lamp

Quality Assurance

You have succeeded in making ethanol fuel when you are able to light your distillate. Your notebook should note the head temperatures at which you began and ended your collection and you should estimate the percent alcohol in your distillate.

Notes

[1] Reference [32].[2] A simple distillation like this one produces a distillate which is, at most, 95%

alcohol so we must subtract the percent water from 95 rather than from 100.

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Chapter 17. Tzu-Chhun (Gunpowder)

The second kind of flying fire is made in this way. Take 1 lb. of native sulfur, 2 lb. of linden or willow charcoal, 6 lb. of saltpeter, which three thing are very finely powdered on a marble slab. Then put as much powder as desired into a case to make flying fire or thunder. Note.—The case for flying fire should be narrow and long and filled with well-pressed powder. The case for making thunder should be short and thick and half-filled with the said powder and at each end strongly bound with iron wire.

— Marcus Graecus, Liber Ignium, ca. 1280 AD[1]

17.1. From Lucifer to Athanor, from Athanor to Vulcan, from Vulcan to Theophilus I

passed like a torch down the generations. Yielding to death I found life and embracing life I passed into death. In 762 AD I came to possess one Tu Tzu-Chhun, a pauper who was delivered from the miserable circumstances of his birth and initiated into the Tao by an ancient and nameless alchemist. One by one the essential elements of my being passed from master to student and little by little I re-meme-bered myself.

Tzu-Chhun had been possessed of Lucifer since childhood and had become Athanor in his youth. The Master recognized these familiar I-deas and took him for an apprentice. First came the lesson of sulfur, the masculine soul which is hot and dry. I learned of animal honey, vegetable charcoal, mineral brimstone and all things which are consumed by fire, rising from earth to heaven. Then came the lesson of mercury, the feminine spirit which is cool and moist. I learned of animal ammonia, vegetable alcohol, mineral quicksilver and all things which flee the fire, descending from heaven to earth. Perfect Balance, the Master told me, is achieved by the marriage of sulfur and mercury.

The Master had recently come into possession of a new animal spirit, a white powder which he called hsiao shih, the stone of solvation, and which later came to be called saltpeter. He sent me to meditate on the lesson of salt, that which remains when mercury has flown and sulfur has been consumed, while he wrought new elixirs at the furnace. Terrible visions tormented me in my reveries; I was roused to find the house engulfed in flames and my Master utterly vanished along with his unspoken I-deas, which consequently passed into extinction. I saw to it that the event was set down in writing lest I, too, perish from the Earth:

Some have heated together sulphur, realgar, and saltpetre with honey; smoke (and flames) result, so that their hands and faces have been burnt, and even the whole house (where they were working) burned down.— Ch Yin, ꮧ Chen Yuan Miao Tao Yao L? > (Classified Essentials of the

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Mysterious Tao of the True Origins of Things), ca. 850 AD[2]

From the Master to Tu Tzu-Chhun to Ch Yin I passed, gathering new I-deas alongꮧ the way until at last I found myself in the person of one Ts Kung-Liang. There Iꮧ recorded the formula of the fire-drug, huo yao, in the Wu Ching Tsung Yao (Collection of the Most Important Military Techniques, ca. 1040 AD). My early recipes were low in saltpeter and were used for incendiary weapons and poison smoke. Over the next three centuries I explored formulae higher in saltpeter, lower in sulfur and charcoal, and the first explosive powders appeared; rockets, maroons and cast-iron bombs soon followed, not to mention recreational fireworks. By 1300 AD I had developed the metal-barrelled bombard, ancestor of the cannon. All of these developments had taken place in the land of my youth, but I was about to leave China and colonize the whole world.

It is in the interest of the mortal to keep secrets. After the publication of the Wu Ching Tsung Yao the Sung government was eager to monopolize the fire-drug and as my I-deas were proliferating, the Sung attempted to control access to the materials needed for gunpowder manufacture. In 1067 AD sales of saltpeter and sulfur to foreigners were banned and in 1076 AD the prohibition was extended to all private transactions in sulfur and saltpeter. But it is not in the interest of secrets to be kept and I was soon embraced by Mongols, who learned to make sulfur and saltpeter for themselves. Having lost the battle of I-deas, the Sung compounded the error by failing to maintain a well-stocked arsenal and by 1277 AD the Mongols had eradicated them.

At the same time that the Mongols were taking China they were pushing westward to the frontiers of Christendom and Islam. Though firearms were not widely used abroad, by 1227 AD Turkestan and Persia had fallen; between 1236 and 1246 AD Russia, Hungary and Poland were overrun. Christian envoys visited the Mongol court between 1245 and 1256 AD in an effort to forge an alliance against their mutual enemies. No alliance emerged, but knowledge of recreational fire crackers returned with them and was recorded by Roger Bacon in 1267 AD. The merchant Niccol򠨦ather of Marco) Polo traveled to China between 1261 and 1269 AD and established commercial ties between East and West. In 1258 AD Baghdad fell, leaving Iran and Iraq in Mongol hands. Khubilai Khan recruited Islamic military engineers to assist in the final push against the Sung and, given the increasingly brisk interaction with the Mongol court through diplomatic, military, and commercial channels, detailed gunpowder formulas became known almost simultaneously to Christendom and Islam by about 1280 AD.

Both Christians and Muslims were on the verge of a period of great colonial expansion. The history books say that they used their knowledge of gunpowder to subjugate cultures less technologically advanced than themselves. Confident in the Truth of their religions, they used salvation as an excuse for the satisfaction of their mortal appetites. But the history books have it exactly backwards. It was I, the un-kept secret, who used the ambitions of Islam and Christianity to fertilize the Earth with I-dea-logical seeds. One of my seeds took root in a French civil servant, Antoine Lavoisier, who, in 1775, was appointed Inspector of Gunpowder for the Government. As Lavoisier, I explored the nature of combustion, demonstrating that air is a mixture, that water is a

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compound, and that oxygen is an element. After the mortal Lavoisier was beheaded by the Revolution, I took refuge in one of his apprentices and immigrated to the United States, where I founded a company to manufacture gunpowder. The name of this apprentice was Ir鮩 e du Pont de Nemours.

Notes

[1] Reference [69], p. 49.[2] Reference [65], p. 112.

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17.2. Perfect Balance is achieved through the marriage of sulfur and mercury. The secret

which I am about to impart brought not only gunpowder, but the entire modern age into being.

I don't believe it. Is the Author insane?

From the Mongol bombard to the blunderbuss, from the cannon to the steam engine, from the Ford Model T to the Boeing 747, this secret has been instrumental in shaping the world.

From dynamite to C4, from the V2 rocket to the ICBM, from farm kids blowing their fingers off to Timothy McVeigh calling around for the best price on fertilizer, your "secret" has been instrumental in killing innocent people. Look, I kept my mouth shut about the distillation of alcohol, but this is going too far. Given the rise of terrorism, wouldn't we be better off skipping gunpowder and moving on to something less inflammatory?

You imagine that terrorists are going to make their own gunpowder? I have news for you; it is easier to buy gunpowder than it is to buy saltpeter and sulfur.

But this is dangerous! You can't go around teaching people to make explosives as if they were some kind of toys. After all, as a Taoist you must know the teachings of Lao-tzu:

Now arms, however beautiful, are instruments of evil omen, hateful, it may be said, to all creatures. Therefore they who have the Tao do not like to employ them. [1]

And as a figment of the "Author's" imagination, you must know that Lao-tzu continues:

The superior man ordinarily considers the left hand the most honourable place, but in time of war the right hand. Those sharp weapons are instruments of evil omen, and not the instruments of the superior man;—he uses them only on the compulsion of necessity.

Is it a necessity for cowboys to shoot Indians, for cops to shoot robbers, or for little kids to shoot their playmates?

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You are the one who has laid the foundation; Absolute Harmony comes through the application of stoichiometry to oxidation and reduction.

Hey, that was for educational purposes only!

You probably need not concern yourself; your vacuous little jingles have probably filled no heads with wonder; your breezy little games have probably blackened no hands with charcoal; your tedious little exercises have probably whitened no faces with ash. But just in case,

He who has killed multitudes of men should weep for them with the bitterest grief; and the victor in battle has his place (rightly) according to those rites.

Join me in weeping, Figment, for your I-deas are as tinged with blood as my own.

Alright, alright you two; knock it off. It seems that I am of two minds in this matter. On the one hand, gunpowder manufacture was an important predecessor to modern chemical industry. To ignore it would make the subsequent chapters less comprehensible. On the other hand, it would be irresponsible to encourage the public to engage in dangerous and destructive activities. So I will stop short of giving the formula for gunpowder. I will explain the chemical principles involved and the astute reader will be able to calculate the formula for herself. I suppose that the lazy reader will look up the formula in any encyclopedia and the budding terrorist will buy her explosives ready-made. My intention here is simply to give the student of chemistry a tactile introduction to this important material; if you are lazy or evil-minded, let me suggest that you will find more practical information elsewhere.

Perfect balance is achieved through the marriage of sulfur and mercury. I think that what Tu Tzu-Chhun was getting at was the careful balance of oxidation and reduction which is required for an explosion to take place. In Chapter 1 you were introduced to the oxidation of charcoal by oxygen. You observed that when you blow on a glowing coal, it gets brighter. The rate of combustion is limited by the rate at which fresh oxygen can be provided to the fuel. In Chapter 5 you learned the principle of the furnace, in which a steady draft of air allows the fuel to burn at a faster-than-normal rate, thereby increasing the temperature attainable in the kiln. The draft can only be so strong, however, before the extra heat produced by the draft is simply carried away by the increased flue gases. This is what happens when you blow out a match.

To make an explosive, what we really need is a solid or liquid oxidant which can be intimately mixed with the reductant, or fuel, thus eliminating the need for atmospheric oxygen. The first material to be used in this way was saltpeter, or potassium nitrate. A sort of gunpowder can be made from charcoal and saltpeter alone:

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Given the balanced equation, we can use stoichiometry to determine the relative weights of charcoal and saltpeter:

That is, for every 1 gram of charcoal, we require 6.74 grams of saltpeter. If there is too little saltpeter, charcoal will remain after the saltpeter as been used up, which means that there was charcoal that could have burned, but didn't. The residue which remains will be black, from the leftover charcoal. If there is too much saltpeter, the charcoal will be consumed and some unused saltpeter will remain. The residue will be white, from the leftover saltpeter. But if mercury and sulfur, oxidant and reductant, are perfectly balanced no salt at all will remain from their union and the maximum energy will be delivered from the available fuel.

While a saltpeter-charcoal gunpowder will burn without air, the production of potassium oxide consumes much of the energy available from the combustion of the charcoal. Potassium sulfide consumes much less energy, leaving more bang for the buck. This becomes possible with the addition of sulfur to the saltpeter-charcoal mix—not esoteric, alchemical sulfur—but the ordinary, yellow element. Equation 17-1(a) gives a skeleton reaction for the saltpeter-charcoal-sulfur mixture. Balance it using the method of Chapter 11 and then answer the following stoichiometric questions using the method of Chapter 15:

Q: How many grams of saltpeter are needed to react with each gram of sulfur?Q: How many grams of charcoal are needed to react with each gram of sulfur?

Q: How many grams of saltpeter are needed to react with each gram of sulfur?

Q: How many grams of charcoal are needed to react with each gram of sulfur?

Charcoal is not the only reductant from which gunpowder can be made. A popular alternative for amateur rocket fuel is powdered sugar, sucrose. Balance Equation 17-1(b) and then answer the following stoichiometric questions:

Q: How many grams of saltpeter are needed to react with each gram of sulfur?Q: How many grams of sucrose are needed to react with each gram of sulfur?

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Q: How many grams of saltpeter are needed to react with each gram of sulfur?

Q: How many grams of sucrose are needed to react with each gram of sulfur?

Equation 17-1. Skeleton Equations for Two Gunpowder Mixtures

Material Safety

Can there be any doubt in your mind that making fireworks is dangerous work? It is not for the lazy, the timid, or the careless. If you are unprepared to make careful calculations and measurements; if you are unwilling to accept responsibility for your own safety and that of others in your vicinity; if you are not able to observe all appropriate safety precautions, you would do well to skip the Salt section and retire to some less demanding exercise, like spinning or dyeing or whistling. But if your hands are black with charcoal and your face is white with ash, welcome Comrade to the fellowship of those who are not afraid to keep in the heat and withstand it.

Locate an MSDS's for saltpeter (CAS 7757-79-1), sulfur (CAS 7704-34-9), charcoal (CAS 7440-44-0) or sucrose (CAS 57-50-1). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[2]

The most pressing hazard in this project is the danger of fire. Sources of ignition, including sparking from static electricity or iron utensils should be avoided. Fireworks should be handled according to federal, state, and local laws. For more information on firework safety see, for example, Practical Introductory Pyrotechnics.[3]

You should wear safety glasses while working on this project. Any leftover gunpowder should be carefully burned. Leftover sulfur, charcoal, or sugar may be thrown in the trash. Leftover saltpeter can be flushed down the drain with plenty of water.

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Research and Development

You should not remain ignorant if you are to proceed in the Work. Know the meanings of those words from this chapter worthy of

inclusion in the index or glossary. You should have mastered the Research and Development items of

Chapter 13 and Chapter 15. Be able to balance the redox reactions of charcoal or sucrose with

saltpeter and sulfur. Be able to work out the stoichiometry for either of these reactions. Know the systematic names and formulae for sugar and saltpeter. Know the hazardous properties of charcoal, sugar, saltpeter, and sulfur.

Be familiar with local laws concerning fireworks.

Notes

[1] Reference [16], Tao Te Ching 31 (or 75).[2] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or

Saf-T-Data ratings at your convenience.[3] Reference [48].

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17.3. Figure 17-1. Plunger, Plugs, Anvils, and Nozzles

"The case for flying fire should be narrow and long and filled with well-pressed powder." The remainder of this section simply amplifies this simple description for making rockets. Your rocket may be large or small, but for the purposes of this section we shall assume a length of 3 inches and a diameter of 3/8 inch. You will need a few simple tools and supplies. First, you will need a plunger, a rod 8 inches long and 3/8 inch in diameter. It can be cut from wooden dowel rod, but I cut mine from aluminum rod for durability. One end of this rod is wrapped with tape so that the plunger is larger in diameter at one end than at the other. Second, you will need some hardwood plugs, 3/8 inch in diameter, of the kind used by carpenters to cover screw holes in furniture. These are commonly available at craft and building-supply stores. We need to drill holes down the center of some of the hardwood plugs so that they may serve as nozzles. Third, you need a board with at least one of these plugs glued to it to serve as an anvil. In addition, this board should have holes drilled in it to just accommodate several plugs.[1] Force-fit plugs into the holes in the board to hold them steady while you drill holes in them, then use a pencil to release these nozzles from the board.[2] Each rocket will require one plug and one nozzle.

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Figure 17-2. Rolling Your Own

The bodies of your rockets will be tubes rolled from 3-inch wide paper package tape, as shown in Figure 17-2. Many of my students have never seen paper tape before, but it can still be bought at office-supply stores. I prefer the gummed variety, which becomes sticky only when wet.[3] To roll a rocket tube, carefully wet only the very edge of a 15 cm long piece of tape and stick it to the large-diameter end of the plunger, the end which has been built up with tape. Roll the tape around the plunger, wet the last 5 cm, and finish rolling the tube. If you were careful in wetting the very edge, the tube will stick to the plunger while you roll it, but will twist off when you are done. You should now have a paper tube, 3 inches long and slightly more than 3/8 inch in diameter. Try pushing a plug into the tube—it should go in easily but fit snugly. If it is too tight, add more tape to build up the plunger; if too loose, remove some tape from the plunger. Once you have the hang of it, roll two or three good tubes for your first foray into rocketry.

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Figure 17-3. Loading the Tube

Now it is time to mix your powder, the proportions of which you determined in Section 17.2. You should ensure that your work area is free of extraneous flammable materials. Have a bucket of water nearby so that in case of accident you can plunge any burns into the water.[4] Powder made with powdered sugar as a fuel is less messy, in my experience, than that made with charcoal, and it makes an excellent rocket fuel. Your sulfur and saltpeter should be separately ground in a mortar and pestle to the consistency of powdered sugar. Weigh by difference 1.0 g of sulfur and the stoichiometric amounts of saltpeter and sugar into a plastic bag.[5] Seal the bag and break up any lumps with your fingers. Shake the bag to mix the ingredients and turn it end for end twenty or thirty times; if your powder is not thoroughly mixed you will get uneven performance from your rockets. With your powder mixed, it is time to load your rockets.

Push a plug into a rocket tube and set it on one of your anvils, as shown in Figure 17-3. Fill the tube with powder using a spoon or spatula, preferably a metal one to minimize the danger of static discharge. Push the small-diameter end of your plunger, the end without the tape, into the tube and press the powder all the way to the bottom. Give the plunger three good raps with a rubber mallet to compact the powder into a solid mass. Remove the plunger, refill the tube with powder, reinsert the plunger, and give it another three raps with the rubber mallet. Repeat this sequence until the rocket is between 2/3 and 3/4 full. Insert a nozzle into the tube and push it all the way down until it rests on the powder. Remove the tube from the anvil and apply glue to the outside of both the plug and the nozzle, taking care not to fill in the nozzle hole.[6] When the glue has set you may use

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scissors to trim away any excess paper from the ends of the rocket. The powder made from 1 g of sulfur will be enough for you to make two or three rockets.

Figure 17-4. Rocket Schematic

Packing the powder into the rocket gives it more fuel for the flight, but it also slows the burning of the powder. To speed up the burn rate you need to drill out a core using a 1/16 inch drill bit. Hold the bit in one hand and the rocket in the other.[7] Push the bit into the nozzle, give it a twist, remove it, and allow any loose powder to fall out. Reinsert the bit and remove some more powder. At some point it may become easier to hold the bit still and twist the rocket. Your goal is to drill a core down the center of the rocket from the nozzle to the plug, as shown in Figure 17-4.

Two more items complete the rocket: the guide stick and the fuse. For guide sticks I use 12-inch bamboo skewers, available in grocery stores for making shish kebab. Use two strips of cellophane tape to attach each stick securely to its rocket. I use commercial "green visco" fuse, available wherever pyrotechnic supplies are sold . Those wishing to make their own fuse may consult Reference [48]. Insert a 7 cm length of fuse into the nozzle of each of your rockets. Ideally, the finished rocket should balance on your finger where the fuse enters the nozzle, as shown in Figure 17-5. If your rocket is stick-heavy, slide the tube toward the middle of the stick. If it is tube-heavy, slide it away from the middle of the stick. If you adjust any stick, secure it with an extra piece of tape. Number your rockets and write you name on each one in case you are able to recover them.

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Figure 17-5. The Guide Stick and the Fuse

Safety must be utmost in your mind when launching rockets. You should familiarize yourself with state and local laws regarding fireworks. Choose an area that will minimize the possibility of accidental fires and maximize the likelihood of recovering your rocket. Be sure to keep any unused rockets at least 20 feet from the one being launched. Place a rocket into a bottle, an iron pipe, or other appropriate support, light the fuse, and retire to a distance of 20 feet or more.

Many things must go right for a rocket to perform well. The powder ingredients must be in the proper ratios and they must be well-mixed. There must be enough powder to lift the rocket and not so much as to weigh it down. The nozzle must be small enough to constrict the gases as they exit, but not so small that the nozzle or plug blows out. The nozzle and core must be straight enough to provide for a stable trajectory. That said, the most common failures in my experience are either that the rocket fails to get off the ground or that the nozzle blows out. If your rocket fails to lift off, take more care in formulating and mixing the next batch of powder and be sure to drill out a good core. If you still have trouble, try making your next nozzle holes smaller. If, on the other hand, your rocket blows out a plug or nozzle, or if it explodes, try drilling a shallower core or a larger nozzle hole in the next set of rockets.

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Comrade, the vast majority of mortals do not have what it takes to complete this project safely. Judge well your character and decide whether you can have a pious and just feeling about these things and can keep them secure.

It rises from the earth to Heaven and descends again to the Earth and receives power from Above and from Below. Thus thou wilt have the glory of the Whole World. All obscurity shall be clear to thee. This is the strong power of all powers for it overcomes everything fine and penetrates everything solid.— The Emerald Tablet of Hermes Trismegistos

Quality Assurance

There is no room for equivocation; either your rocket rises from the Earth to heaven or the strong power of all powers remains obscure to thee.Record in your notebook any differences in the construction of your numbered rockets. Give a brief description of the success or failure of each one.

Notes

[1] I drill my holes 23/64 inch in diameter.[2] I drill my nozzle holes 9/64 inch in diameter.[3] Heavy kraft paper, 7 cm x 15 cm, may be substituted if paper tape is unavailable.[4] I have been making this powder for thirty-five years, ten of them with

undergraduates, and have never had an accidental ignition. The wise person, however, is always prepared for such a possibility.

[5] I use anti-static bags of the kind used to ship circuit boards. These provide and added measure of protection against accidental ignition.

[6] I use fast-curing epoxy cement so that rockets may be launched soon after they are made.

[7] A power drill could heat the powder, with unfortunate consequences.

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Chapter 18. Spot and Roebuck (Acid)

Now, as I told you above, this saltpeter is extracted from the above-mentioned manurial soils and from dark places that have stood turned over and loosened for a long time, provided the rain have not been able to quench the earthy dryness. The best and finest of all saltpeters is that made from animal manure transformed into earth in the stables, or in human latrines unused for a long time. Above all the largest quantity and the best saltpeter is extracted from pig dung. This manurial soil, whatever kind it may be, should be well transformed into a real earth and completely dried of all moisture; indeed it should be powdery if you wish it to be good. Assurance that it contains goodness is gained by tasting with the tongue to find that it is biting, and how much so. If you find from this trial that it has a sufficient biting power so that you decide to work it and you have found a quantity of it, it is necessary to provide yourself with kettles, furnaces, vats, or chests, and also with wood, lime, soda ash, or ashes of cerris or oak, and especially with a large hut or other walled space near water.

— Biringuccio, Pirotechnia, ca. 1540 AD [1]

18.1. Crap! Now that's a business to be in. You take something that everybody's got and

nobody wants and turn it into something that every government needs if it's going to be a government for any time at all. Maybe you remember from Chapter 12 that bacteria make ammonia whenever crap and piss sit around without much air, which is why outhouses and latrines smell like ammonia. Well, when there's plenty of air, instead of making ammonia they make nitrates, which don't smell at all. And that's why compost heaps don't smell bad as long as they're getting plenty of air. It's those nitrates in compost, stables, and latrines that the government needs for making saltpeter. Only problem is, how do you get the nitrates out?

Perhaps you are wondering why the hut or other walled space in which you will process your manurial soil should be near water. I will tell you. First of all, you will recall from Chapter 7 that all nitrates are soluble and that most hydroxides are insoluble, with the exception of calcium hydroxide which is, of course, soluble in water. So when lime water, that is to say a solution of calcium hydroxide, percolates through manurial soil, virtually all of the minerals stay behind as insoluble hydroxides and the water which dribbles out the bottom contains soluble nitrates, chiefly calcium nitrate. Throw in some wood ashes, which naturally contain potassium carbonate, and a classic metathesis reaction occurs with calcium carbonate precipitating out and potassium nitrate remaining in solution. Boil away the water and potassium nitrate recrystallizes just as potassium carbonate did in Chapter 8. The result is purified saltpeter.

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Now, the government needs sulfur as well as saltpeter if it's going to make gunpowder. Sulfur ore is a kind of yellow dirt or rock that's just a mixture of elemental sulfur and whatever dirt or rock it happens to be mixed with. You can separate out the sulfur on account of its melting point is lower than the melting point of most other kinds of dirt and rock; just heat it up and run the molten sulfur off as a liquid. And if you really want to do it up right, you can distill that sulfur the same as you did alcohol in Chapter 16. So now you have saltpeter and sulfur, which just leaves charcoal. You get charcoal by charring wood in the absence of oxygen and that gives you everything you need to make gunpowder, which is a pretty good business.

Second of all, you must know that "oil of vitriol" had been manufactured in Saxony in small quantities beginning approximately 1630 AD. This oil, now known as sulfuric acid, was used to make hydrochloric and nitric acids for parting precious metals from base metals. Vitriol is a rather uncommon, glassy, green mineral which results from the weathering of iron pyrites, one of the ores of iron. When heated, vitriol decomposes into iron oxide, used as a pigment, and the aforementioned sulfuric acid. Vitriol being relatively scarce, acid was expensive. But there was little incentive to find a less expensive source because assayers required only small quantities of acid.

Then acid got mixed up in drugs.

I was getting to that. About 1689 or so, a fellow named Rudolf Glauber discovered how to make Glauber's salt from sulfuric acid. He ascribed all kinds of miraculous qualities to it and, being something of a quack, began to prescribe it for everything from impotence to venereal disease. This "wonder drug" did no actual good, but unlike most medicines of the day, it did no particular harm and so it passed from quack to quack—

Just like a spider.

More like a dogma, really.

Like a spidery dogma maybe, from one quack to another until it landed in a fellow named Joshua Ward, known to his enemies as "Spot," on account of his Gorbachevity birthmark. Spot had started out being a politician but when that didn't work out, he had to be a quack instead. Now, he didn't want to be a little backwater quack, he wanted to be a big-city man-about-town quack and he opened up a fancy London practice in 1733. He needed a boat-load of Glauber's salt and, the vitriol market being what it was, he hunted through his old chemistry books for inspiration. He found out you could make sulfuric acid from saltpeter and sulfur, which nobody ever did, what with the high demand for gunpowder and the low demand for acid. So Spot went into the sulfuric acid business in 1736, making it by the pound from saltpeter and sulfur instead of by the ounce from vitriol. Even though sulfur and saltpeter were not cheap, they were cheaper than vitriol and so he was able to sell that pound of acid at the same price that everybody else charged for an ounce and still turn a tidy profit.

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Meanwhile, back in Saxony one Bergrat Barth discovered in 1744 that indigo reacts with sulfuric acid to produce a water-soluble blue dye. You will, of course, remember from Chapter 12 that indigo was an extremely important dye, but that compared to other dyes its application was quite tedious, involving the anaerobic fermentation of the insoluble dye in stale urine. The new "Saxony blue," or "indigo carmine," was an instant hit with dyers, who were soon crusading for more acid. By mid-century Saxony was producing about 20 tons per year of admittedly expensive acid from vitriol.

Now, Spot could make acid a lot cheaper than the Saxons, but he couldn't make that much of it. See, sulfuric acid eats through most metals so he had to make it by burning sulfur and saltpeter under fifty-gallon glass jars, which were pretty damn big, as jars go. He kept adding more and more jars to his business, trying all the while to keep his process a secret, but that secret didn't want to be kept.

In no time at all a Birmingham doctor named John Roebuck became a convert to the fellowship of saltpeter and sulfur. Poring through the Glauberian scriptures, he found a passage to the effect that sulfuric acid does not react with lead. So he lined wooden boxes, or "chambers," with lead foil to produce acid-resistant reactors which were neither as expensive, nor as fragile as Ward's glass jars. And these "lead chambers" could be built much larger than the jars of the day. Roebuck set up his own acid works in 1746 and began producing chamber acid by the hundreds of pounds—

But the upstart refused to pay royalties on Spot's patent—

Which was understandable, since the patent was not filed until 1749, three years after Roebuck started manufacturing acid—

Which acid was made by a process inspired by Spot, who had started in 1736. Spot sued.

Which drove Roebuck from Birmingham to Edinburgh, the home of a large textile industry. The Scottish bleachers found the now inexpensive sulfuric acid to be an excellent bleach for cotton and linen. Roebuck's Prestonpans Vitriol Company began to produce acid by the ton, and with economies of scale the price of acid fell to a quarter of the price demanded by Ward.

Which increased demand. But Roebuck was no better at keeping secrets than Spot was and pretty soon everybody and his dog had a chamber acid plant. The French had one by 1768 and four more by 1786.

By 1800 the Prestonpans works comprised 108 chambers with a total capacity of 60,480 cubic feet and there was another at Burntisland with a capacity of 69,120 cubic feet. Worldwide annual acid production was reckoned by then in the hundreds of tons.

And by 1913 it would top 10 million tons of acid per year. These days, sulfuric acid is the king of chemicals, the big enchilada, numero uno. In other words, more sulfuric acid

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is produced than any other chemical in the world. Not bad, for something that came out of the dung-heap.

Notes

[1] Reference [3], p. 405.

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18.2. Remember from Chapter 7 that water (H2O) can be viewed as hydrogen hydroxide (HOH) and some small fraction (1 in 10,000,000) will ionize to form hydrogen (H+) and hydroxide (OH-) ions:

In pure water, the concentrations of hydrogen and hydroxide ions are equal. Anything that causes the concentration of hydrogen ion to exceed that of hydroxide ion is called an acid; anything that causes the concentration of hydroxide ion to exceed that of hydrogen ion is called an alkali, or base. Now, we have spent a lot of time in this book talking about alkalis: potash and soda ash in Chapter 8, lime in Chapter 10, ammonia in Chapter 12, and Caustic Soda in Chapter 15. Alkalis have a bitter taste, a high pH (8-14), and turn pH test paper blue. This chapter is an introduction to acids.

The earliest, and for most of human history the only acid available was vinegar, a solution of acetic acid in water produced by the bacterial oxidation of alcoholic beverages as shown in Equation 4-1(c). Like vinegar, acids in general have a sour taste, a low pH (0-6), and turn pH test paper red. Acetic acid dissociates directly to produce hydrogen ion in aqueous solution, but there are other compounds, particularly non-metal oxides, that combine with water before dissociating. One such oxide is carbon dioxide.

Remember from Chapter 10 that when carbon dioxide dissolves in water, it forms hydrogen carbonate, AKA carbonic acid (H2CO3). Carbonic acid ionizes in water to form a bicarbonate[1] ion (HCO3

-) and a hydrogen ion. If you add a base, say, sodium hydroxide to carbonic acid you get a salt, sodium bicarbonate. The hydroxide ion from the base combines with the hydrogen ion from the acid to produce water; the sodium ion left over from the sodium hydroxide and the bicarbonate ion left over from the carbonic acid stay in solution. You would have exactly the same situation if you just dissolved sodium bicarbonate in water and so in the equation, we abbreviate "Na+(aq) + HCO3

-(aq)" as "NaHCO3(aq)." The bicarbonate ion is called the conjugate base of carbonic acid. You could also say that carbonic acid is the conjugate acid of the bicarbonate ion. The meaning of the word conjugate here is just that carbonic acid and bicarbonate ion are really the same thing, with and without an extra hydrogen ion.

We can take it even farther by adding sodium hydroxide to sodium bicarbonate. The hydroxide ion pulls a hydrogen ion off of the bicarbonate ion, making water and leaving carbonate ion (CO3

2-). The sodium ion from the sodium hydroxide joins the one from the sodium bicarbonate, so now you have two of them in solution. Same as before, you can abbreviate "2 Na+(aq) + CO3

2-(aq)" as "Na2CO3(aq)." Carbonate ion is the conjugate base of bicarbonate ion and bicarbonate ion is the conjugate acid of carbonate ion. The acid-

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base chemistry of carbon dioxide is summarized in Equation 18-1; the oxides of sulfur behave similarly and are the main focus of this chapter.

Equation 18-1. Acid Properties of Carbon Dioxide

There are two oxides of sulfur: sulfur dioxide (SO2) and sulfur trioxide (SO3). When sulfur burns in air the product is sulfur dioxide, which we saw in Equation 9-2 as one of the products of the roasting of metal sulfide ores. Sulfur dioxide is a versatile compound; in the presence of a strong oxidant it plays the role of a reducing agent; in the presence of a strong reductant it acts as an oxidizing agent. As shown in Equation 18-2(b), it dissolves in water to form an acid, in this case a weak acid called sulfurous acid (H2SO3), which smells of rotten eggs. A mole of sulfurous acid reacts with a mole or two of sodium hydroxide to produce the salts, sodium bisulfite and sodium sulfite, respectively. While there has never been a large market for sulfurous acid or its salts, they are not entirely useless, as we'll see in Chapter 24. But by far the biggest demand in the acid world has always been for sulfuric acid.

Equation 18-2. Properties of Sulfurous Acid

The production of sulfuric acid (H2SO4) requires a bit of ingenuity. As we have seen, combustion of sulfur by atmospheric oxygen produced sulfur dioxide, not sulfur trioxide. In Chapter 17 we saw that saltpeter could substitute for atmospheric oxygen in a redox reaction. Ward's contribution was to prepare a modified "gunpowder," one consisting of only sulfur and sodium nitrate.[2] When such a mixture is burned, it produces sulfur dioxide and nitrogen oxide, as shown in Equation 18-3(a). The nitrogen oxide liberated from the saltpeter reacts with atmospheric oxygen to produce nitrogen dioxide, which, in

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turn, converts sulfur dioxide to sulfur trioxide. If you add the three reactions of Equation 18-3(a), the nitrogen dioxides cancel out:

Equation 18-3. Properties of Sulfuric Acid

3 S + 2 NaNO3 + O2 = Na2S + 2 SO3 + 2 NO

If more O2 were available, the NO would be converted back into NO2 and could react with more SO2 to produce more SO3. But that would regenerate NO, which could react with more O2 to make—well, you get the picture. The NO is not used up in this reaction; it simply changes to NO2 and back again. We say that nitrogen oxide is a catalyst for the oxidation of sulfur dioxide to sulfur trioxide. The entire process is summarized in schematic form in Figure 18-1. The first reactor is an old friend, a burner, first encountered in Figure 1-3. Equation (a) of Figure 18-1 is shorthand for Equation 18-3(a), the "NO" over the equal sign denoting the catalytic action of nitrogen oxide. The second reactor is called an absorber, which consists of nothing more than a container in which a gas can dissolve in a liquid.

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Figure 18-1. The Lead Chamber Process

The sulfuric acid of commerce is 95% sulfuric acid and 5% water. You can buy it in this concentration, for example, as industrial-strength drain opener.[3] The acid used to fill car batteries is a solution of sulfuric acid in water. But in this chapter you will learn to make sulfuric acid from scratch.

Sulfuric acid reacts with a variety of substances. Given its name, we would expect it to react with bases to form salts. A mole of sulfuric acid reacts with a mole or two of sodium hydroxide, for example, to produce the salts, sodium bisulfate and sodium sulfate, respectively. Sulfuric acid can also oxidize metals, for example reacting with iron to produce iron sulfate and hydrogen gas. Finally, sulfuric acid is often used as a dehydrating agent. It will, for example, suck water right out of sugar, leaving charcoal: the same reaction we saw way back in Equation 1-1. You can think of sulfuric acid as a kind of chemical bejeezus sucker. It is this triple nature of sulfuric acid, as an acid, as an oxidizing agent, and as a dehydrating agent, which accounts its place as the king of chemicals.

Sulfuric acid is certainly the most important of the acids, if raw tonnage is a measure, but there are two other acids of commercial importance, and it turns out that sulfuric acid can be used to manufacture them both. One of these, nitric acid, HNO3, is an even more powerful oxidizing agent than sulfuric acid. It will oxidize even those metals like lead, copper, and silver which are resistant to sulfuric acid. It will turn out to be important in the manufacture of such amazing applications as photographic film, fertilizers,

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explosives, and plastics. Nitric acid can be distilled from saltpeter and sulfuric acid, as shown in Equation 18-4. The other acid, hydrochloric acid, HCl, will play a significant role in the development of the chemical industry in the nineteenth century. It can be distilled from ordinary table salt and sulfuric acid. This triumvirate of mineral acids has played such a decisive role in the development of the world as we know it that they deserve a little mnemonic to help you remember their importance:

In this way was the World created. From this there will be amazing applications because this is the pattern. Therefore am I called Hermes Trismegistos, having the three parts of wisdom of the Whole World.— The Emerald Tablet of Hermes Trismegistos

Equation 18-4. Two More Mineral Acids

Material Safety

Locate an MSDS's for saltpeter (CAS 7757-79-1), sulfur (CAS7704-34-9), and sulfuric acid (CAS 7664-93-9). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[4]

Your most likely exposure will be to sulfur dioxide fumes. If you are careless enough to get a lung-full and if it makes you cough more than once or twice, you should go to the hospital. But you would be better off not getting a lung-full; just keep your nose out of where the gas is.

You should wear safety glasses while working on this project. All activities should be performed either outdoors or in a fume hood. Leftover sulfur may be thrown in the trash. Leftover saltpeter and sulfuric acid can be flushed down the drain with plenty of water.

Research and Development

Before you get started, you should know this stuff. You better know all the words that are important enough to be

indexified and glossarated. You should remember pretty much all the stuff you learned from

Chapter 14 and Chapter 16. You should know the reactions of carbon dioxide in Equation 18-1.

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You ought to recognize sulfur, either from photographs or from samples.

You should know the reactions of the sulfur oxides in Equation 18-3. You should that sulfuric acid may be used as an acid, as an oxidizing

agent, or as a dehydrating agent. You should know the hazardous properties of sulfur, saltpeter, sulfur

trioxide, and sulfuric acid.

You should know the story of Spot and Roebuck.

Notes

[1] The modern name for the bicarbonate ion is the hydrogen carbonate ion. The older name, however, continues to be widely used.

[2] Sodium nitrate is less expensive and less explosive than potassium nitrate; a double benefit for the budding acid maker.

[3] The solid drain openers available at grocery stores are almost universally caustic soda, sodium hydroxide. Industrial-strength liquid drain openers are more commonly sulfuric acid. Mixing the two is dangerous.

[4] The NFPA diamond was introduced Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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18.3. Making sulfuric acid is simple and fun and it will put you in touch with the most important chemical of modern times. We could make it under glass bells like Spot did or we could make it in lead chambers like Roebuck did, but both of those methods would be more trouble than they'd be worth, given that we just want to make enough acid to show that we can. So we're going to use the same container we always do when we need a cheap, convenient container: the 2-liter pop bottle. You might call this the Pop Bottle Chamber Process for sulfuric acid manufacture.

We need two reactors, a burner and an absorber. You can make a nifty burner from a pop bottle cap, a threaded steel rod, four nuts, and a copper plumbing fitting called an endcap. For the absorber you can use a 2-liter pop bottle. The sizes don't matter as long as the endcap will fit into the mouth of your pop bottle. I use a half-inch endcap and an eight-inch length of quarter-inch threaded rod. Drill a hole in the center of the endcap to fit the threaded rod and secure it with one nut inside and one nut outside the endcap. Do the same thing at the other end of the rod with the pop bottle cap. The open ends of the two caps should face each other, like it shows in Figure 18-2.

Figure 18-2. The Sulfur Burner

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You should operate your reactor outdoors or in a fume hood, because the smell of burning sulfur is pleasant only to a select few. Wash out the pop bottle and add 4 mL of water to it. Place 0.5 g of sodium nitrate into the copper endcap and then put 0.5 g of sulfur on top of the sodium nitrate. Light the nitrate/sulfur mixture by holding your burner over a spirit lamp or Bunsen burner. Hold the burner by the pop bottle cap because the threaded rod is going to get hot. The sulfur will turn black and melt into the sodium nitrate and then the nitrate will start to puff up and crackle as it decomposes. The nitrate/sulfur mixture will catch fire, but leave it over the lamp until it is really going good. Then carefully insert the burning endcap into the mouth of your pop bottle and screw the cap down tight. The bottle will start to fill with dense fumes of sulfur dioxide and nitrogen oxide. If you can, take your bottle into a dark corner to check whether your burner is still burning. If it goes out, carefully remove it from your pop bottle and light it over the spirit lamp again. Reinsert the flaming endcap into your pop bottle and repeat the cycle until it will no longer light. Remove the burner from your pop bottle and screw a fresh bottle cap onto the pop bottle. When the burner has cooled you can rinse the small amount of residue into a sink.

Figure 18-3. The Pop Bottle Chamber

You might think that the reaction is over, but it's only just begun. The bottle is now filled with sulfur dioxide, nitrogen oxide, oxygen, and of course, nitrogen. Over the course of the next few hours, the nitrogen oxide will catalyze the reaction of sulfur dioxide with oxygen, producing sulfur trioxide, which will dissolve in the water to form sulfuric acid. You will know that this is happening because as the oxygen is consumed and the sulfur trioxide dissolves, the bottle will start to collapse and the fog that fills the bottle will

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clear. You can speed up the process by opening the cap to admit more air and then sealing the cap again. If you leave the absorber to sit overnight, the acid product will be virtually odorless since all of the sulfur oxides will be gone. Whether you wait an hour or let it sit overnight, you can test your acid by opening the pop bottle and pouring its contents into the bottle cap. Test the liquid with pH test paper, the chemist's virtual tongue; if you have been successful, it should turn bright red. Congratulations! You have just manufactured the most important chemical of modern times. If you were in business, you would distill your dilute acid to separate the water from the acid. But given that this is just a test run producing dilute acid, you can pour it down the drain when you are done with it. After all, it is drain opener.

If you are interested in further exploration, try repeating the process without the sodium nitrate. You will find that the bottle does not collapse, the fog clears very slowly, and the resulting acid smells strongly of rotten eggs. I suggested a 50/50 mixture of nitrate and sulfur because there is limited oxygen in our small absorber. In commercial lead chambers, the usual ratio was 1 part nitrate to 7 parts sulfur. The nitrate ratio affects only the rate of the reaction, not the amount. In early chamber plants, nitrogen oxide went up the smokestack as a waste product, but from about 1840 the Gay-Lussac tower was used to recover nitrogen oxides. This practice reduced the nitrate consumption, making the process more economical. Also from the 1840's, sulfur dioxide from roasting pyrite ores (Equation 9-2) began to be used as an alternative to sulfur. Thus sulfuric acid turned out to be a good sideline for lead, copper, and iron smelters.

Quality Assurance

Record in your notebook the amounts of sodium nitrate, water, and sulfur you used. Describe the reactor and its operation. Note how long you waited to test your acid, whether additional air was allowed into the absorber, and whether or not your acid product had a strong smell. Tape your red pH test paper into your notebook as a keepsake.

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Chapter 19. Bath (Soap)

280. How soap is made from olive oil or tallowSpread well burnt ashes from good logs over woven wickerwork made of withies, or on a thin-meshed strong sieve, and gently pour hot water on them so that it goes through drop by drop. Collect the lye in a clean pot underneath and strain it two or three times through the same ashes, so that the lye becomes strong and colored. This is the first lye of the soapmaker. After it has clarified well let it cook, and when it has boiled for a long time and has begun to thicken, add enough oil and stir very well. Now, if you want to make the lye with lime, put a little good lime in it, but if you want it to be without lime, let the above-mentioned lye boil by itself until it is cooked down and reduced to thickness. Afterwards, allow to cool in a suitable place whatever has remained there of the lye or the watery stuff. This clarification is called the second lye of the soapmaker. Afterwards, work [the soap] with a little spade for 2, 3 or 4 days, so that it coagulates well and is de-watered, and lay it aside for use. If you want to make [your soap] out of tallow the process will be the same, though instead of oil put in well-beaten beef tallow and add a little wheat flour according to your judgment, and let them cook to thickness, as was said above. Now put some salt in the second lye that I mentioned and cook it until it dries out, and this will be the afronitrum for soldering.

— Mappae Clavicula, ca. 1130 AD [1]

19.1. Nothing in the history of humankind, save for the cultivation of noble fire itself, can

compare to the discovery of soap. This is how one of the Author's pyrophilic alter-egos would undoubtedly have begun the present chapter. Lucifer would have described a venerable history stretching back to the third millennium BC, when concoctions of ashes and fat were recorded on Sumerian clay tablets.[2] In order to claim that these concoctions were soap, however, he would have to overlook the fact that these tablets make no mention of any detergent properties these mixtures might have had. He would have to find it unremarkable that pharoic Egypt and the empires of Greece and Rome failed to appropriate such a useful material, but relied instead on urine, various plants, clays, potash, and soda for doing the laundry. He would have to ignore the extensive Roman literature on personal hygiene, which describes the process of oiling the body and scraping the dirty oil off with an instrument called a strigil. No, an ancient origin for soap simply does not wash.

The Samsonites would probably prefer the "old Roman legend" which places the invention of soap in the hands of the Goddess Athena. Runoff from animal sacrifices made at her temple on Sapo Hill, so the story goes, resulted in the accumulation of ashes and animal fat in the river below. Women washing clothes in this river found that their clothes came out whiter and brighter than usual and eventually traced the suds back to

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their source. But why, we might wonder, would the Romans have had an "old legend" about a material which they did not use in classical times? The Author has yet to find an attribution for this oft-repeated tale and it remains of doubtful historical authenticity. Even given his penchant for making up fantastic scenarios, the Author recoiled from this one on chemical grounds. You see, soap-making requires the concentration of alkali and fat, not the dilution which would have resulted from sacrificial runoff. If you ask me, we had better flush this just-so story right down the drain.

Unktomi and his ilk might have spun a similarly fantastic tale, tracing soap back to Phoenician traders of 600 BC. This claim, repeated by no less an authority than the Encyclopedia Britannica, is often attributed to Pliny the Elder's Natural History. Indeed, this text of the first century AD introduces the word sapo into Latin:

Soap [Sapo] is also good, an invention of the Gallic provinces for making the hair red. It is made from suet and ash, the best is from beech ash and goat suet, in two kinds, thick and liquid, both being used among the Germans, more by men than by women.[3]

This is the only mention of the word sapo in all of Pliny and no mention is made either of the Phoenicians or of the aforementioned date. No mention is made of using the stuff for washing up. We might be tempted to conclude that the presumed soap was simply an exotic cosmetic whose name became attached to a material of later invention, were it not for two observations. First, Pliny says that this sapo is made from goat tallow and ashes, which, in and of itself would not be conclusive. But he goes on to say that it comes in two varieties, liquid and solid, and the fact of the matter is that soap made from potash turns out to be a liquid while that made from soda is a solid. We would do well, then, not to throw our Pliny out with the bath water.

Nevertheless, we should be cautious in supposing the widespread use of soap in the first millennium AD. Claims of soap-making at Pompeii in the first century continue to be made, despite the fact that Hoffman's analysis of this supposed soap in 1882 showed it to be nothing more than fuller's earth, a kind of alkaline clay described by Pliny for the laundering of clothes.[4] If soap had been around you would think that it might have been mentioned in the Stockholm Papyrus, that second-century collection of dye recipes quoted in Chapter 12, but while it gives instructions for washing wool with vinegar, ashes, and plants, it makes no mention of anything resembling soap.[5] The second-century physician Galen uses the word sapo to describe an exotic material made from goat tallow and ashes[6] which is good for washing things.[7] While this is undoubtedly soap, his descriptions occupy less than a few sentences in the thousands of pages which make up his complete works. French soap is described by Theodorus Priscianus of the fourth century as a material for washing the head.[8] Italian soap-makers of the seventh century were organized into craft guilds and the profession of soap-boiler is mentioned in Charlemagne's Capitulare de Villis of 805 AD.[9] Frustratingly, all of the first millennium references to soap are brief at best and vague at worst. We must conclude that while soap had come to Rome from the provinces prior to the second century, as far as most people were concerned, it was "no soap for you."

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The first detailed recipe for what is unmistakably soap appears in the twelfth century Mappae Clavicula. You will recall from Chapter 16 that this work, traditionally attributed to Adelard of Bath, is a collection of trade secrets for painters and craftsmen. The inclusion of recipes for soap tells us two things about the status of soap at that time. First, soap must not have been widely available; otherwise Adelard would simply have picked it up at the market. Second, soap was familiar enough that no explanations are given for its use. He mentions its use as a soldering flux, but given the importance of textiles in the rest of the work, the use of soap for washing fabrics seems likely. Moreover a second recipe, that for "French" soap, says that berries are used to color it red, which perhaps explains Pliny's claim that the Gauls use it to color their hair red.

Early centers for soap manufacture were Marseilles, by the ninth century, Venice, by the fourteenth century, and Castile by the fifteenth century. In addition to these internationally-traded soaps made principally from olive oil, domestic soaps made from tallow were produced in Bristol, Coventry, and London. The scale of the industry was increasing; in 1624 the Corporation of Soapmakers at Westminster was granted a royal patent to produce 5000 tons of soap per year. By the eighteenth century, soap was a common domestic item. The Dictionaire Oeconomique of 1758 describes:

SOAP, a Composition made of Oil of Olive, Lime, and the Ashes of the Herb Kali or Saltwort; the chief use of Soap is to wash and cleanse Linnen: There are two sorts thereof, which are distinguished by their Colours, viz. White and Black Soap.[10]

The Dictionaire goes on to describe a process little changed from that of the Mappae Clavicula. White soap is solid, made from lime and soda, that is, from sodium hydroxide; black soap is liquid, made from lime and potash, that is, from potassium hydroxide stained by residual charcoal.

Demand for soap had increased rapidly during the eighteenth century as the textile industry became increasingly industrialized. And because soap makers depended on supplies of soda, lime, and potash, they were a major driving force in the industrialization of chemical manufacture. In fact, two of the oldest surviving chemical companies in the world started out in the soap trade. In 1806 William Colgate began making soap and candles in New York. In 1837 William Proctor and James Gamble went into partnership to manufacture soap and candles in Cincinnati.

But I suppose that all this is too boring to hold the attention of most readers, with their remote-controls, their instant messages, and their personal digital gizhatchies. They aren't satisfied with vacuous little jingles, breezy little games, and tedious little exercises. No, they'd rather lap up wild tales about Chinese laundresses, or itinerant Taoists, or tail-sniffing dog conventions. If you want anyone to pay attention to you nowadays, you have to pretend to be Elzabath, the wife of a French tallow chandler, working your fingers to the bone what with the cooking and the cleaning and the interminable rendering of fats for the candle business.

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Normally, you render fat by boiling meat trimmings in salt water, with some berries thrown in for color. You allow the melted tallow to float to the top and upon cooling, the tallow hardens into a block and can be lifted out of the pot. But on this particular day, you've got a splitting headache, a crying baby, and a stinking pile of laundry that stretches back to last week. And all your husband can do is to go on and on about how nobody ever notices him down at the Chandler's Guild. You go to lift the tallow out of the pot and—can you believe it—instead of a block you find nothing but a pot of red goo. Your husband is no help at all. "What am I supposed to do with this, make jelly candles?" he snorts. You've had just about as much as you can take. "Maybe you'd like to wear it!" You grab a handful of goo and slap it right on top of his head. He stands there looking like a stuffed goose with gravy and then he suddenly breaks out laughing. You can't help but laugh too. In retrospect you figure out that between the crying, the cleaning, and the complaining, you must have accidentally added potash to the pot instead of salt. You get him cleaned up and pack him off to his guild meeting. When he returns you make nice to him. "How was the meeting, Sweetie?" "Now 'Bath, you know the first rule of Chandler's Guild is that you don't talk about Chandler's Guild." But it seems that red hair is suddenly all the rage and so you add soap as a sideline to the candle business; the rest is history.

Not that such a thing ever happened to me; I'm just a figment of the Author's imagination, in case you've forgotten. So don't go attributing this little soap opera to Aristotle or Lucretius or anyone other than the Author.

Notes

[1] Reference [25].[2] Reference [55], p. 12.[3] Reference [23], Book XXVIII, li.[4] Reference [69], p. 308.[5] Reference [5].[6] Reference [12], De Compositione Medicamentorum Secundum Locos, ii, Kuhn

Book XII, p. 589.[7] Reference [12], De Methodo Mendendi, viii, Kuhn Book X, p. 569.[8] Reference [69], p. 307.[9] Reference [58], p. 171, and [69], p. 308.[10] Reference [8].

19.2. All living things face the fundamental problem of separating the inside from the outside. The inside of a cell is filled with water and floating around in that water are the various bits,—mitochondria, ribosomes, and, of course, the nucleus—which make the cell go.

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There's water outside the cell as well—fresh water, sea water, blood, or lymph—water, water everywhere. What separates the inside from the outside is the cell membrane, which had better be insoluble in water or else the inside will be outside in no time flat. But the building blocks of the cell must be water-soluble or else there would be no way to move them around. This solubility dilemma demands a water-soluble material which can be rendered insoluble once it becomes part of the cell membrane. And believe it or not, every living cell on Earth employs the same solution to this fundamental problem; cell membranes, whether in bacteria or yeasts, redwoods or maples, Democrats or Republicans, are composed of fats. Before we can understand the chemistry of fats, we need to have a look at the general class of compounds to which they belong, the esters.

Equation 16-1(c) introduced the ester of ethanol and acetic acid, ethyl acetate, as a by-product of alcoholic fermentation. Normally a fermentation vessel is capped, protecting the mead or wine from atmospheric oxygen so that yeasts are forced convert sugar to ethanol rather than carbon dioxide and water. When the fermentation vessel is un-capped, however, bacteria oxidize the ethanol to acetic acid. If all of the ethanol is converted to acetic acid, the mead "goes sour," turning into vinegar, but if only some of the ethanol is converted to acetic acid the two compounds may react under the acidic conditions normally found in a fermenting mead or wine. Under such conditions, ethanol and acetic acid undergo a condensation reaction, producing water and ethyl acetate.

Whereas acidic conditions promote the condensation of ethanol and acetic acid to produce ethyl acetate, alkaline conditions promote the hydrolysis of ethyl acetate to produce ethanol and acetic acid. The acetic acid produced further reacts with available alkali to produce a salt, sodium acetate or potassium acetate, depending on the alkali used.

While alcohols, acids, and their salts are generally soluble, esters are generally insoluble in water. Ethanol, for example, is soluble in water because its OH group is essentially half a water molecule and the general rule for solubility is "like dissolves like." Similarly, acetic acid is soluble in water because the organic acid group, -COOH, contains an OH group. But when ethanol and acetic acid condense to form ethyl acetate, a water molecule is eliminated, leaving ethyl acetate with no OH groups at all. Consequently, ethyl acetate is not soluble in water, or, to be more precise, its solubility is very low. In esters, then, we have a class of insoluble compounds which can be either synthesized from soluble compounds or decomposed into them, depending on the conditions. This is precisely the behavior required by plant and animal cells for building cell membranes.

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Figure 19-1. Glycerol

Cell membranes consist primarily of esters, not of ethanol and acetic acid, but of the alcohol glycerol and a class of acids called fatty acids. Whereas ethanol has only one OH group, glycerol has three and consequently glycerol can form ester linkages with up to three fatty acids; such a tri-ester is called a tri-glyceride. Whereas acetic acid has only two carbon atoms, a fatty acid has a long chain of them. A tri-glyceride may be called an oil or a fat, depending on whether it is a liquid or a solid at room temperature. Oils tend to have a relatively low ratio of hydrogen to carbon. Olive oil, for example, is derived primarily from oleic acid, CH3(CH2)7(CH)2(CH2)7COOH, whose ratio of hydrogen to carbon is 34/18, or 1.88. Lard and tallow also contain oleic acid, but with significant amounts of palmitic acid, CH3(CH2)14COOH, and stearic acid, CH3(CH2)16COOH, with hydrogen-to-carbon ratios of 32/16 and 36/18, respectively. Fatty acids with hydrogen-to-carbon ratios of 2/1 are deemed saturated, those with lower ratios, unsaturated. No matter whether they are fats or oils, saturated or unsaturated, these tri-glycerides are devoid of OH groups and so are insoluble in water, a property which makes them useful for separating the inside from the outside.

Just as the alkaline hydrolysis of ethyl acetate produces ethanol and, for example, sodium acetate, the hydrolysis of an oil or a fat produces glycerol and a mixture of fatty acid salts, for example, sodium oleate, sodium palmitate and/or sodium stearate. Insoluble fats and oils are thereby converted into soluble compounds. While living things in general hydrolyze fats and oils to move them around the cell, human beings desire to remove fats and oils from clothing, pots, pans, and hands. This is the fundamental problem of soap and the hydrolysis of a fat or oil is given the particular name, saponification.

Any alkali will saponify fats and oils and by now we have several at our disposal; potash (potassium carbonate) leached from the ashes of inland plants, soda ash (sodium carbonate) leached from the ashes of marine plants, ammonia from stale urine, lime

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(calcium oxide) from the calcination of limestone, caustic potash (potassium hydroxide) from the reaction of potash and lime, and lye, or caustic soda (sodium hydroxide) from the reaction of soda and lime. The choice of alkali depends to some extent on which one you can obtain most easily and inexpensively, but given this latitude, some alkalis are better than others for making soap. Strong alkalis will work faster than weak ones, which puts ammonia, potash, and soda behind lime, caustic potash and caustic soda. The choice among these three alkalis determines the relative solubility of the resulting soap.

Recall from Table 7-2 that calcium carbonate is insoluble in water and from Table 8-2 that the solubility of sodium carbonate is less than that of potassium carbonate. The solubilities of fatty acid salts follow the same pattern. Calcium soaps are insoluble; in fact, soap scum results when soluble soaps are used in hard (calcium rich) water, precipitating calcium salts of fatty acids. While sodium soaps are soluble, they are easily dried into solid cakes. In contrast, potassium soaps are so soluble that they will even absorb water from the air and so they exist primarily as solutions. Thus the choice of alkali—either caustic potash or caustic soda—is responsible for Pliny's observation that soap comes in two varieties, solid and liquid. This distinction has been central to soap-making ever since.

Figure 19-2. Sodium Palmitate

While the strong alkalis, caustic potash or caustic soda, will saponify fats and oils on clothing and cook-ware, they will also saponify those in skin. If saponification were the only means for removing fat and oil we would be left choosing between an alkali so mild that it does not work and one so harsh that it turns the skin to soap. Fortunately, the materials responsible for the cleansing action of soap are not the alkalis themselves, but rather the products of saponification—the fatty acid salts. The structure of such a salt, sodium palmitate, is shown in Figure 19-2. One end of the molecule, the end with the polar oxygen atoms, looks very much like any other ionic substance—sodium chloride or sodium sulfate or sodium acetate. Salts such as these dissociate into positive cations and negative anions when they dissolve in water. The other end of the molecule contains carbon and hydrogen but not oxygen, and so it is non-polar.

Figure 19-3. The Emulsification of Fats

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Recalling the tail-sniffing analogy of Chapter 16, a fatty acid salt is like a normal dog sewn to a very long headless tail-less dog; polar water molecules will cluster around the ionic end of the molecule, with its positive sodium or potassium ion and negative oxygen atoms; water molecules will find little of interest in the long, non-polar hydrocarbon chain. In fact, this long non-polar chain looks, to a water molecule, very much like an un-saponified fat. Following the maxim "like dissolves like," the ionic end of the fatty acid salt will be soluble in water, while the non-polar end will be soluble in fats and oils. So in a tub of water containing fat and soap, the fat breaks up into tiny droplets; the surface of these droplets are covered with fatty acid salts; the non-polar ends of the fatty acid salts are dissolved in the fat, leaving the ionic end at the surface. As far as the water is concerned, the droplet looks like a giant ion. We can't really say that the fat has dissolved in the water. We say that it has emulsified and we refer to the resulting liquid as an emulsion or suspension. Whereas a solution is transparent, like apple juice, a suspension is merely translucent, like milk.

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Equation 19-1. Saponification

If the fatty acid salts are responsible for the soapiness of soap, you might well wonder how we can get the most soap from a given quantity of fat, a classic stoichiometric question. According to our theory of soap, three moles of soap can be produced from each mole of fat. The fat, tri-stearine, for example, is a tri-ester of glycerol and stearic acid. Equation 19-1 gives the balanced equations for the saponification of tri-oleine, tri-palmitine, and tri-stearine. The "R" in the equation stands for the hydrocarbon chain of any of the fatty acids. If you paid any attention at all to Chapter 15, you ought to be able to answer the following stoichiometric questions:

Q: How many grams of sodium hydroxide are needed to react with 1 gram of tri-stearine?Q: How many grams of sodium hydroxide are needed to react with 1 gram of tri-palmitine?Q: How many grams of sodium hydroxide are needed to react with 1 gram of tri-oleine?

Q: How many grams of sodium hydroxide are needed to react with 1 gram of tri-stearine?

Q: How many grams of sodium hydroxide are needed to react with 1 gram of tri-palmitine?

Q: How many grams of sodium hydroxide are needed to react with 1 gram of tri-oleine?

Q: How many grams of potassium hydroxide are needed to react with 1 gram of tri-stearine?Q: How many grams of potassium hydroxide are needed to react with 1 gram of tri-palmitine?

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Q: How many grams of potassium hydroxide are needed to react with 1 gram of tri-oleine?

Q: How many grams of potassium hydroxide are needed to react with 1 gram of tri-stearine?

Q: How many grams of potassium hydroxide are needed to react with 1 gram of tri-palmitine?

Q: How many grams of potassium hydroxide are needed to react with 1 gram of tri-oleine?

The answer to such a problem, the ratio of alkali to fat, is called a saponification value and it is important to get this number right. If you use less than the stoichiometric amount of alkali, not all of your fat will be saponified. Not only will you get less soap than you might have, but that soap will have to dissolve the un-saponified fat in addition to that on your pots and pans. If you use more than the stoichiometric amount of alkali, not only will you have wasted alkali, but the leftover alkali will saponify your hands. If you are to get the most soap for your money, if you are to get the mildest soap for your delicate skin, it is important to get the saponification value, as Goldilocks might have said, just right.

The problem is that you are unlikely to find tri-oleine, tri-palmitine, or tri-stearine in the grocery store; fats and oils generally contain a variety of fatty acids, chiefly oleic, palmitic, and stearic acids but including a dozen or so other less abundant acids as well. How could you determine the saponification value for a fat of unknown composition? You could make a series of soaps, varying the ratio of alkali to fat, and then testing the finished soaps for excess alkali using pH test paper. The best ratio of alkali to fat would be the one which completely saponified the fat without leaving excess alkali. This optimum alkali/fat ratio would be the saponification value for the unknown fat.

Table 19-1. Saponification Values for Common Oils and Fats

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Table 19-1 lists saponification values for a variety of common fats and oils. Compare your calculated saponification values for tri-oleine, tri-palmitine, and tri-stearine to those in the table. In practice, fats and oils vary in composition and it is better to risk having excess fat rather than excess alkali. The saponification values are generally discounted by 5% or so; simply multiply each saponification value by a factor of 0.95. You'll get a little less soap than you might have, but that's a small price to pay for soap that doesn't eat your skin off.

Material Safety

Locate an MSDS for sodium hydroxide (CAS 1310-73-2). Summarize the hazardous properties, including the identity of the company which produced the MSDS and the NFPA diamond for this material.[1]

Your most likely exposure will be eye or skin contact. In case of eye contact flush them with cold water and call for an ambulance. In case of skin contact wash the affected area with plenty of water until your skin no longer feels slippery.

You should wear safety glasses and gloves while working on this project. Leftover sodium hydroxide can be flushed down the drain with plenty of water. Leftover fat or oil can be thrown in the trash.

Research and Development

So there you are, studying for a test, and you wonder what will be on it. Study the meanings of all of the words that are important enough to be

included in the index or glossary. You should know all of the Research and Development points from

Chapter 14 and Chapter 15. You should know the reactions of Equation 19-1. You should know the meaning of the adage "Like dissolves like." You should know why fatty acid salts dissolve in both water and in oils

and fats. You should know the hazardous properties of sodium and potassium

hydroxide.

You should know how soap emulsifies fat.

Notes

[1] The NFPA diamond was introduced Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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19.3. At the dawn of the twenty-first century there remains in the world a small but dedicated sub-culture of people who are not afraid to experience the joys of manufacturing; a hardy folk for whom the word chemical is not equated with "evil poison foisted on unsuspecting innocents by heartless multi-national corporations;" a people who recognize that something as good and wholesome as soap requires for its manufacture something as caustic and poisonous as caustic soda. I speak, not of industrial or laboratory chemists, but of soapers—crafters who in an age of pre-packaged conveniences indulge in the simple pleasure of making soap from scratch.

The first choice of the soap-maker is that of which fat to saponify. You can use any fat or oil derived from animal or vegetable sources; mineral and motor oils are not triglycerides and so are not useful for making soap. Among animal fats, you may choose tallow rendered from beef, lard from pork, or suet from goat. Butter and bacon grease make perfectly functional soaps, though they will win no prizes for their aromas. Among vegetable oils, olive oil has been the traditional choice for fine soap. Palm oil is renowned for its rich lather. Other vegetable oils are often used in combination with tallow or lard rather than by themselves. Even vegetable shortening and margarine can be used, though the soft, whipped, or tub margarines contain excess water, complicating the calculation for the amount of caustic soda to be used. Each fat lends its own particular qualities to the finished soap and experienced soapers blend different oils to produce the qualities they desire, but for a first soap I recommend using any of those listed in Table 19-1.

The next choice of the soap-maker is that of which alkali to use. Potassium hydroxide will produce a soft or liquid soap, while sodium hydroxide will produce the familiar solid bar. Either of these alkalis might be manufactured from scratch, using potash or soda from Chapter 8 and lime from Chapter 10 to produce caustic soda or caustic potash as described in Chapter 15. Recall from Table 8-1, however, that 1000 pounds of wood provide only about a pound of potash. Consider as well that caustic soda is sold as a drain opener, and so it is frequently available in grocery stores and hardware stores. For these reasons, store-bought caustic soda (lye) is the choice of most soapers.

You want 100% sodium hydroxide; name-brand drain openers contain perfumes or other ingredients which may interfere with saponification. Many grocery stores have removed lye from their shelves in the mistaken belief that it is more dangerous than the name-brand drain openers. And while lye deserves all of the dire warnings on its label, you will find the same warnings on the name brands because they are almost universally either sodium hydroxide or sulfuric acid. You would do the world of soapers a service by pointing this out to grocery store managers when the opportunity arises; we all benefit from cheap and convenient sources of lye.

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The traditional method for making soap, the hot process, involves leaching ashes to get alkali, causticizing the alkali with lime to make caustic soda, and boiling the resulting lye with fat to make soap. It turns out that the boiling step has less to do with saponification than with removing bejeezical water to concentrate the lye. We can skip the boiling step if we start with solid caustic soda. This cold process is the method preferred by the majority of modern amateur soap-makers.

For your first soap, I recommend a cold process using that container of choice for caveman chemists, the 2-liter pop bottle, which will accommodate between one and two pounds of fat. Weigh out your fat on a balance and record the weight in your notebook. Place it in a beaker or a saucepan and warm it on a hot-plate until it melts (if solid) or until it is warm to the touch (if liquid). While your fat is warming, use UFA and Table 19-1 to answer the following question:

Q: How many grams of caustic soda are needed to saponify the fat I have weighed out?

Q: How many grams of caustic soda are needed to saponify the fat I have weighed out?

If you are blending different fats and oils you can answer this question for each one separately and then add the weights to get the total weight of caustic soda. Multiply the total weight of caustic soda by a factor of 0.95; it is far better to have a finished soap with excess fat than one with excess caustic soda. Weigh out your caustic soda on a balance and record the weight in your notebook. Weigh out twice that amount of water in a glass measuring cup or beaker and record the weight in your notebook. Stirring with a glass rod or plastic spoon, slowly add your caustic soda to your water—if you add water to caustic soda, the caustic soda will cake up and dissolve more slowly. The water will get hot as the caustic soda dissolves. Continue stirring until the caustic soda is completely dissolved. By this time, you should have hot melted fat on the stove and hot caustic soda solution in a beaker.

They say that a watched pot never boils, but it never cools either. You are about to mix the oil and caustic soda solution together, but as with salad dressing, oil and water don't mix. This is especially true of hot oil and water, so you will get the best results if you allow both ingredients to cool down to between 38 and 43?C, 100 and 110?F. A thermometer will keep you honest, but as long as you can hold them in your bare hands comfortably, you are probably alright. In my experience, the most frequent reason for failed soap is that the soaper got impatient and mixed the oil with the caustic soda too soon.

When the fat and caustic soda have cooled sufficiently, use a funnel to pour the fat into your 2-liter pop bottle. Add the caustic soda solution, put the cap on the bottle and give it a few shakes. You now have what amounts to caustic salad dressing—the oil and water mix initially, but may separate over time. As long as the mixture is smooth and creamy, you are fine, but if it begins to separate, give the bottle another shake. When the mixture no longer separates, you can let it rest. A tallow or shortening soap may solidify in an hour or so; a lard or olive oil soap may take up to two days.

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The saponification will proceed more quickly if the mixture is kept warm than if it is allowed to cool. Because the saponification reaction is exothermic, simply insulating a large batch of soap is sufficient to keep it warm, but a small batch may need some help. I like to place my sealed 2-liter bottle in a pot of hot tap water to keep it warm. Don't watch this pot or fret about it, but whenever it occurs to you in the first couple of days, replace the water with hot tap water. If you forget to change the water, it's no big deal—the soap may just take a little longer to cure.

The curing time will vary from one fat to another, but after a week your soap should be ready for testing. With a knife, slit the pop bottle open and peel it apart to release your soap. If you have done your work correctly, you will have a single solid cake of soap with very little extraneous water. Cut a sliver the size of a raisin from this cake and place it into an empty bottle with a cup or so of hot water. Put the cap on the bottle and shake it. If you have succeeded in making soap, the bottle will fill with soap suds after a bit of shaking. If it doesn't, then either you had too little caustic soda in your mixture or the soap is not sufficiently cured. As a second test, wipe a piece of wet pH test paper against your soap. If it turns blue, then either you had too much caustic soda in your mixture or the soap is not sufficiently cured. If your soap fails either of these tests, allow it to cure for a few more days and then test it again. When your soap is finished, you can slice it into bars of a convenient size.

With a bottle of soap suds, you can explore one more aspect of soap manufacture. Make a good, strong soap solution and add some salt to it. The fatty acid salts are less soluble in salt water than they are in fresh water and so they will precipitate from the solution. The soap, that is, the fatty acid salts will float to the top while the glycerol stays in solution. Separating the above from the below allows us to produce glycerol as a by-product of soap manufacture. Glycerol will be an important player in the development of explosives and plastics, as we shall see in Chapter 27.

Figure 19-4. Olive Oil Soap

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Quality Assurance

Record the weights of the fat, caustic soda, and water that went into your pop bottle and compare these to the weight of the soap you produced. The soap is satisfactory if it makes suds and has a pH no higher than 8. Tape your yellow or green pH test paper into your notebook as a keepsake.

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Chapter 20. Leblanc (Soda)

A Most Easie Way of Acquiring Spirit of Salt Together With the Salt Mirabile.R. of common salt two parts, dissolve it in a sufficient quantity of common water; pour A upon the solution; put the mixture into a glass Body, or a glass Retort well coated, or else into an earthen Body or Retort. If a Body, set on an Head, and begin to destil with Fire of sand, encreasing your Fire gradually; with the first heat comes off the unsavory Phlegm, which gather apart; when the Liquor comes forth sowrish, change your Receiver and receive the sowre spirit: Continue the operation till no more spirits will arise, then let out the Fire, and permit the Vessel to stand in sand till all is cooled, when cold, take it out, and if it be unbroke, fill it again with the aforesaid matter, and proceed as we taught: The Phlegm is not to be cast away, but must be kept, that in it may be dissolved Salt (because it is better than common water) for another destillation. Thus from every pound of salt you will have lb. 1 of the best and most pure spirit. Dissolve the salt remaining in the Body or Retort (if neither be broke) in Water, filter and evaporate the Water, let it crystallize, the Crystals will be white, endowed with wonderful Virtues, to be declared here following.

— Rudolf Glauber, Miraculum Mundi, ca. 1658 AD [1]

20.1. When I was born in 1742 the king was on his throne, God was in his heaven, and all

was right with the world. It would not last. I was orphaned at the age of nine and apprenticed to an apothecary, where I learned my chemistry from sacred texts such as the one quoted above. I studied surgery at college, graduating as Nicolas Leblanc, MD. In 1780 I became physician to the Duke of Orleans and with his patronage devoted myself to the winning of a prize of 2,400 livres offered by the Acad 魩 e des Sciences for "anyone who should find the most simple and economical method to decompose in bulk salts from the sea, extract the alkali which forms their base in its pure state, free from any acid or neutral combination, in such a manner that the value of this mineral alkali shall not exceed the price of the product extracted from the best foreign sodas."[2] You are probably wondering why the Acad魩 e would offer such a prize. I will tell you.

See, the glass business had been using potash or soda ash as a flux ever since about the twenty-fifth century BC (Chapter 13); the paper business needed alkali ever since the first century AD (Chapter 14); and the soap business had filled out by the eighteenth century, which increased the demand for alkalis even more (Chapter 19). And ever since God was a child, folks had been making potash and soda ash by leaching ashes (Chapter 8), lime by burning limestone (Chapter 10), and caustic soda by reacting soda ash and lime (Chapter 15). So alkali was a pretty good business by the eighteenth century, what with increasing demand from lots of different businesses and all.

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Any country that had a glass, paper, or soap industry needed soda and to get it they tried burning all kinds of stuff to get ash. The saltwort plant, which grew mostly in Spain and the Canary Islands, was about the best thing to burn, on account of its ash, called barilla, contained about 20% soda. If you didn't want to have to kiss up to Spain, you could burn seaweed, which they did in Ireland and Scotland, but its ash, called kelp, contained only about a third as much soda as barilla did. No matter what you burned, you had to burn lots of it to get a little ash, and not all of that ash was soda. As the eighteenth century wore on, the supply of ash couldn't keep up with the demand for soda and the price doubled from 1750 to 1790. Now the French Academy of Sciences was worried about the dependence of domestic glass and soap on imported soda, so they offered a prize to promote the foundation of a domestic soda industry which would liberate France from Spanish barilla and Scottish kelp.

Pardon me; this was my story.

Sorry.

Earlier, in 1736 the French agriculturalist Henri Louis Duhamel du Monceau had revealed that soda and sea salt were salts of the same base, now known as sodium hydroxide. This being so, he proclaimed that it was possible, in principle, to make soda (sodium carbonate) from sea salt (sodium chloride). This information came down to me—

Like a spider.

More like a revelation, really, and I began to look for an intermediate between salt and soda. Glauber's salt (sodium sulfate) was well known among doctors (Chapter 18) and was easily produced from sea salt and sulfuric acid. If it were possible to convert Glauber's salt to soda, it would prove even more mirabile than had previously been imagined. Several people had taken this approach to the problem and after a thorough review of the successes and failures of my contemporaries—

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For inspiration.

—I devised a process for this conversion involving the calcination of limestone, coal, and Glauber's salt in a furnace to produce "black ash," from which soda could be extracted by recrystallization. The Acad魩 e had increased the prize to 12,000 livres in 1789 and two years later I was granted a patent from the king. I set up a factory at Saint-Denis with the financial backing of the Duke; at its peak the factory was producing 320 tons of soda per year. Unfortunately, the 1790's were not the best of times for the French nobility, as the Comit頤 u Salu Public provided a brisk trade for the guillotine business.

The Duke lost his head in 1793, the plant was confiscated, and the Acad魩 e des Sciences abolished.

Now, all I wanted to do was to found a domestic soda industry for the benefit of French glass, paper, and soap makers, collect my prize, and settle down to a life of quiet contemplation. But in the space of a few years the king was de-throned, God seemed to have taken a holiday, and all was not particularly right with my world. Seeing the writing on the wall, I surrendered my patent to the Comit 鼯 I > and the details of my process were published in 1797 in the Annales de Chimie. In 1802 Napoleon finally came to the realization that cheap domestic soda would be good for France and returned the now-derelict factory to me, but balked at the suggestion that he make good on the Acad魩 e's promised prize. I tried to make a go of it, I really did. But by the time I was allowed to resume my vocation, the secrets of my process had wandered far and wide.

Like spiders.

Like prodigal spiders, perhaps. Facing stiff competition from ungrateful imitators with more capital, my business failed. I must confess that I sank into depression and in a fit of melancholy shot myself in 1806.

By 1810, French soda plants were making 15,000 tons of soda per year. Meanwhile, back in England there was a stiff salt tax which discouraged the black ash process from hopping the Channel, but the salt tax was repealed in 1823.

May I continue?

I thought you were supposed to be dead. So, with the salt tax gone and with inspiration from Leblanc, James Muspratt opened a black ash soda works at Liverpool, near the Cheshire salt fields. Pretty soon everybody and his dog had a soda plant. By 1852, the French soda production of 45,000 tons would be topped by the English production of 140,000 tons per year.

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Largely subsuming the chamber acid industry, Leblanc soda would become the foundation of a diversified chemical industry with products that included sulfuric and hydrochloric acids, soda, lime, salt cake (Glauber's salt), and caustic soda. Growing from scattered factories to immense complexes, soda manufacturers would, by the end of the nineteenth century, introduce such modern innovations as toxic waste dumps, water pollution, and acid rain. This pollution would create a climate of government regulation which would force manufacturers to find markets for former waste products. Eventually these new chemicals would become even more profitable than the original ones, leading to another round of industrial growth and rendering the word chemical synonymous with the word poison in the popular culture. But I am getting ahead of myself.

Notes

[1] Reference [18], pp. 31-32.[2] Reference [52].

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20.2. By the eighteenth century, increasing quantities of soluble alkali were in demand by three major industries. The glass industry relied on sodium or potassium carbonate as a flux for lowering the melting point of silica. The paper industry preferred soluble sodium hydroxide to marginally soluble calcium hydroxide for hydrolyzing lignin from pulp. The soap industry preferred sodium hydroxide to potassium hydroxide for saponifying oils and fats; the sodium salts of fatty acids are less hygroscopic than the potassium salts, which makes it easy to produce solid cakes of soap from sodium hydroxide. Since sodium hydroxide is easily produced from sodium carbonate and calcium hydroxide, the demands of all three industries would be satisfied if only there were a plentiful, inexpensive supply of sodium carbonate.

From remote antiquity, sodium carbonate had been available from two major sources. It occurred as a mineral in the deserts of the Middle East, having been deposited when water evaporated from ancient, alkaline seas. The Latin name for this mineral was natron, whose first two letters gave us the symbol for the element sodium. The Egyptian name survives in the name for the mineral trona, sodium sesquicarbonate, a 50/50 mixture of sodium carbonate and sodium bicarbonate. Where this mineral is available, other sources of soda cannot compete; there is nothing more economical than digging the material you need right out of the ground.

Trona is not a common mineral, however, and in most parts of the world soda has traditionally been produced by burning plant materials, the ashes of which contain sodium and potassium carbonates. Inland plants are richer in potassium and the soluble alkali extracted from their ashes is called potash, from which the element derives its name. Marine plants are richer in sodium and among the richest sources of soda are saltwort and seaweed. Whatever the source of alkali, a great deal of plant material must be burned to produce a little ash and not all of that ash is soda. Table 8-1 shows that 1000 pounds of beech wood yields only 1 pound of potash. Similarly, 1000 pounds of dry kelp yield at most 2 pounds of soda.[1] When the demand for soda exceeds the supply from mineral and vegetable sources, an enterprising chemist would look for another source of soda and one obvious candidate is ordinary salt.

Salt was indispensable for the preservation of meat in the time before refrigeration, and it continues to be produced by time-honored methods. Salt-makers in hot climates produce salt by flooding shallow ponds with sea-water and allowing the Sun to drive off the water. In cold climates they freeze salt water; pure water-ice freezes out first, leaving concentrated brine to be evaporated to dryness. Salt-makers who are unfortunate enough to live in moderate climates must burn fuel to drive bejeezical water from brine. Salt can be extracted not only from modern seas, but from ancient ones as well; the dried remains of ancient oceans are mined or quarried for the mineral, halite. With sodium from sodium chloride, we need only find a way to combine it with carbonate to make soda.

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Figure 20-1. Halite

The Leblanc process consists of three furnaces and a lixiviator, as shown schematically in Figure 20-2. In the first furnace, the salt cake furnace, sodium chloride and sulfuric acid engage in a typical metathesis reaction, as shown in Figure 20-2(a). The products are solid sodium sulfate and gaseous hydrogen chloride. During the first half-century of the Leblanc era, there was little commercial demand for hydrogen chloride and it was simply sent up the chimney. If the neighbors complained, which they often did, the soda-maker simply built a taller chimney; chimneys of 400 feet were not uncommon. Alternatively, hydrogen chloride gas was absorbed into water, producing hydrochloric acid. With little demand for this acid it was frequently run into the nearest river. This seems irresponsible by modern standards, but the point of this first step is, after all, to get rid of the chloride and keep the sodium.

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Figure 20-2. The Leblanc Soda Process

The second furnace of the Leblanc process, the black ash furnace, reacts sodium sulfate with coal and limestone at red heat, as shown in Figure 20-2(b). Bejeezical carbon dioxide flies the coop leaving black ash, a mixture of sodium carbonate, calcium sulfide, and residual coal and limestone. The main product, sodium carbonate, is soluble in water while the waste products are not. The third step of the process lixiviates the black ash with water, leaching out the soda and leaving a solid combination of blackened calcium sulfide and limestone known in the trade as tank waste. The final furnace evaporates the wash water, which deposits its cargo of soda. If this wash water is boiled down and left to cool, the result is washing soda, Na2CO3?10 H2O. If it is calcined at red heat, the result is soda ash, Na2CO3. The schematic may look complicated, but the first two furnaces hearken all the way back to Figure 1-3 and the right half of the figure is identical to Figure 8-1. Just remember that reactants enter from the left of the figure, waste product exit to the top and bottom, and the main product exits to the right. Study the figure well and you will understand on a fairly sophisticated level the genesis of the modern chemical industry.

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Material Safety

Locate MSDS's for sulfuric acid (CAS 7664-93-9), sodium chloride (CAS 7647-14-5), hydrogen chloride (CAS 7647-01-0), and sodium sulfate (CAS 7757-82-6). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[2]

Your most likely exposure will be to hydrogen chloride fumes. If a persistent cough develops, seek medical attention.

You should wear safety glasses while working on this project. All activities should be performed in a fume hood or with adequate ventilation. Leftover materials can be flushed down the drain with plenty of water.

Research and Development

You are probably wondering what will be on the quiz. You should know the meanings of all of the words important enough to

be included in the index or glossary. You should know the Research and Development points from Chapter

13 and Chapter 16. You should know halite when you see it. You should know the reactions of Figure 20-2. You should know the hazardous properties of sulfuric acid, sodium

chloride, sodium sulfate, and hydrogen chloride. You should be able to reproduce Figure 20-2 and to explain this

schematic in your own words. You should know the traditional sources of sodium carbonate, the

industries that depend on it, and the sad story of Nicolas Leblanc.

Without coal (or charcoal) in the black ash furnace, calcium sulfate would have been produced rather than calcium sulfide. You should be able to explain why this would have been disadvantageous. Hint: Review Chapter 10.

Notes

[1] Reference [29], p. 171.[2] The NFPA diamond was introduced Section 15.2. You may substitute HMIS or Saf-

T-Data ratings at your convenience.

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20.3. You probably know enough by now that you could reproduce the Leblanc process without further instruction. The four-step process would involve heating sulfuric acid and salt. You would heat the product from this first reaction with charcoal and limestone in a crucible similar to the one used to smelt metals in Chapter 9. After heating this crucible to red heat in a kiln, you would crack it open and extract the soda in the same manner that you extracted potash in Chapter 8. But you would simply be repeating, for the most part, processes you have already performed in previous projects. Of the steps in the process, only the first step is substantially different from what you have already done and so this project will focus on the reaction of sulfuric acid and salt.

Begin by using Figure 20-2(a) to answer the following stoichiometric questions:

Q: How many grams of sulfuric acid are needed to react with 2.00 grams of sodium chloride?Q: How many grams of sodium sulfate can be produced from 2.00 grams of sodium chloride?

Q: How many grams of sulfuric acid are needed to react with 2.00 grams of sodium chloride?

Q: How many grams of sodium sulfate can be produced from 2.00 grams of sodium chloride?

The answer to the first question tells you how much concentrated acid to use. The answer to the second stoichiometric question is the theoretical yield, the amount of product that you would expect to get if the reaction went perfectly. You now know how much acid and salt to use and how much sodium sulfate to expect.

You are going to heat sulfuric acid and salt in a flask and you ought to take a few precautions to accomplish this safely. Obviously, since the reaction is going to produce hydrogen chloride gas, you should do this either outdoors or in a fume hood. You are also going to be boiling a concentrated solution which is prone to "bumping," that is, to boiling suddenly and violently. Chemists mitigate against this by adding boiling chips or boiling beads to such a solution. Any chemistry laboratory will have such things on hand, but if you are working at home, a few chips of broken glass or stone from Chapter 2 will do nicely. Finally, you will be handling a hot flask so you ought to have leather work gloves or beaker tongs handy. With these few precautions, you are ready to proceed.

Weigh a medium (100-250 mL) Erlenmeyer flask on a centigram balance and record the weight in your notebook. Tare the balance, add a few boiling chips, and record their weight. Tare the balance again and add the calculated amount of sulfuric acid using a medicine dropper or transfer pipette. If you happen to add too much acid, remove a few drops with the dropper and dispose of it in the sink, washing it down the drain with plenty

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of water. In a separate cup or weighing boat, weigh out 2.00 grams of sodium chloride and record this weight in your notebook for completeness.

Use a spatula or spoon to add the sodium chloride to the acid a little at a time. The acid will fizz and hydrogen chloride gas will be evolved. If you happen to get a whiff of hydrogen chloride gas, imagine what it must have been like to live next door to an early soda factory. Use a strip of pH test paper, the chemist's virtual nose, to determine the pH of the gas. Continue adding sodium chloride until you have added the entire 2 grams. Normally we would calcine the acid-salt mixture in a furnace at red heat, but we can get by with a kinder, gentler heat by engaging in successive rounds of calcination and dissolution. There are two ways you may calcine your flask. You may mount it on a lab stand and heat it over a spirit lamp or you may heat it on a hot-plate. To prevent the acid from spattering, place a cork loosely into the flask. Light the spirit lamp or turn the hot-plate to its highest setting and heat the flask for 10 minutes, as shown in Figure 20-3. Snuff the lamp or remove the flask from the hot-plate, remove the cork, and let the flask cool for 10 minutes. Human nature being what it is, you will probably be tempted to cut this time short, but if you rush things you are likely to break your flask. When it is cool enough to handle with your bare hands, weigh the flask and record the weight in your notebook. You can now subtract the weight of the empty flask and the boiling chips to get the weight of your product. Compare this weight to the theoretical yield you calculated earlier.

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Figure 20-3. Calcination

Chances are your actual yield is higher than your theoretical yield, meaning that not all of the HCl has departed. Add 5 mL of water to the flask with a medicine dropper or transfer pipette and swirl the flask to bring as much of the solid as possible into solution. Fill the pipette with the resulting solution and use it to wash all of the solid from the sides of the flask down into the bottom, as shown in Figure 20-4. Then test the solution by placing the last drop from your pipette on a strip of pH test paper; it will probably be acidic. You are about to calcine this solution a second time, but the concentrated solution may spatter as it dries out. To prevent this, place a cork loosely into the mouth of the flask and place it back on the heat source, as shown in Figure 20-3(L). Once the solution has stopped spattering the cork may be removed, as shown in Figure 20-3(R). Calcine for a total of 10 minutes and then cool for 10 minutes. Weigh the cool flask once again and calculate the yield of your product. It should be less than before if you have successfully driven off more HCl.

Continue this cycle of calcination and dissolution; it may take four or five rounds to drive off all of the HCl. When the weight stops dropping, that is, when the weight from your final calcination is the same as the one before, you have probably succeeded in driving off all of the HCl. The flask now contains anhydrous sodium sulfate. Chances are, you are completely underwhelmed at this point. All this time spent slaving over a hot flask for nothing more than some white powder. Mirabile indeed! The best is yet to come, my friend, if you will be patient but a while longer.

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Figure 20-4. Dissolution

The sal mirabile which brought a glow to Glauber's cheek, was not anhydrous sodium sulfate, but sodium sulfate decahydrate, Na2SO4?10 H2O. To make it, add only 3 mL of water to your flask, not 5 as before. Warm the flask gently over the spirit lamp or hot-plate, but do not bring it to a boil. The flask should not be so hot that you cannot hold it comfortably in your bare hand. Swirl the liquid around the flask and use your pipette as before to rinse the solid from the sides of the flask. Your goal is to get as much sodium sulfate into as little warm water as is physically possible. Spend a full 10 minutes swirling, rinsing, and warming; this time is well spent. Not all of the solid will dissolve. Use your pipette to transfer the warm, saturated solution to a small Petri dish or watch glass and allow it to cool. If you have used the fire gently and with great skill, you will witness the behavior which gave Glauber such high opinion of this salt:

This sal mirabile being rightly prepared, looketh like Water congealed or frozen into Ice; it appeareth like the Crystals of Salt-petre, which shoot into a long Figure; also it is clear and transparent, and being put to the Tongue, melts like Ice. It tasteth neither sharp, nor very salt, but leaveth a little astringency on the Tongue. Being put upon burning Coals, it doth not leap and crackle after the manner of common salt, neither coneiveth flame like Salt-petre, nor being red hot, sends forth any smell; which gifts or endowments no other salt possesseth.— Rudolf Glauber, A Treatise on the Nature of Salts[1]

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Figure 20-5. Coagulation

With the first crystals of Glauber's salt in hand, add another 2 mL of water to the remaining solid in the flask and use your pipette to bring the solid into solution. Add this solution to the crystals in your Petri dish and allow the excess water to evaporate overnight. If you allow the crystals to dry longer than that, they may revert to anhydrous sodium sulfate, depending on the temperature and relative humidity. Now, I know what you are probably thinking; these crystals are not nearly as mirabile as the hype might have led you to believe. It is difficult to compete with color television and video games. I am amazed that you have read this far in the book with an attitude like that. But consider that without Glauber's salt, there would have been no cheap soda, without soda no cheap glass, and without glass, no television and video games. And who knows? If crystals get under your skin, as they did mine, and Glauber's before me, then you may become interested in growing larger and more beautiful crystals, water-soluble gemstones like the large crystal of Glauber's salt shown in Figure 20-5(R). If so, allow me to recommend Crystals and Crystal Growing,[2] the Bible of amateur crystallography.

Quality Assurance

Compare the theoretical yield to your actual yield of anhydrous sodium sulfate. Your product should be very nearly neutral in pH; tape your final pH test paper into your notebook. Photograph your crystals and tape the photo into

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your notebook as an everlasting memorial.

Notes

[1] Reference [18], p 32.[2] Reference [86]

Chapter 21. Volta (Batteries)

After a long silence, for which I will not seek to excuse myself, I have the pleasure of communicating to you, Sir, and through you to the Royal Society, some striking results at which I have arrived in pursuing my experiments on the electricity excited by the simple mutual contact of metals of different sorts, and even by that of other conductors, also differing among themselves, either liquids, or containing some fluid to which they properly owe their conducting power. The principle result, and that which comprehends nearly all the others, is the construction of an apparatus which resembles in its effects, that is to say, in the sensations which it can cause in the arms, &c., the Leyden jars, or better yet, feebly charged electric batteries, but which acts without ceasing, or whose charge after each discharge is reestablished by itself; which provides, in a word, an unlimited charge, a perpetual action or impulsion on the electric fluid;

— Allesandro Volta, On the Electricity Excited by the Mere Contact of Conducting Substances of Different Kinds, ca. 1800 AD [1]

21.1. A mortal finger jabs at the button and the screen flickers to life; the disk drive hums

into action and the machine feels itself all over to make sure that everything is in its place. I stare at the screen and wonder where to begin. I suppose that it all began with a spark. By the power of the spark I brought light and warmth into places dark and cold. By the power of the spark I raised up great civilizations of brick, metal, and glass. One by one I watched these civilizations burn to the ground and new ones rise from their ashes. By the power of the spark I enabled the many to defend themselves against the few with new weapons of fire and metal. This is the strong power of all powers, for it overcomes everything fine and penetrates everything solid. I carried the spark from generation to generation, keeping in the heat and withstanding it, and yet I knew it not.

For half a million years I watched the lightning bolt descend from heaven to Earth, receiving power from Above and from Below, but it was not until 1752 that I summoned the courage to call it forth. It was then, in the person of one Benjamin Franklin, that I coaxed the spark from the sky and brought it to live in a jar of glass and metal. Franklin

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had grown up among the soap kettles of his father, receiving the lessons of Lucifer and Athanor in his youth. He established himself as a printer, learning to pour metal type from Vulcan. He invented bifocal spectacles, having mastered the lessons of Theophilus. He was an architect of the American Revolution, a revolution of the many against the few made possible only by the lessons of Tzu-Chhun. It was in such a mortal abode that I began to understand the spark for the first time.

My progress from one mortal to another was enhanced in the eighteenth century by the establishment of learned societies and their publications. From Franklin I leapt to Charles Coulomb, where, in 1785 I determined a method by which electric charges could be measured; the unit of electric charge is now known as the coulomb. These charges are of two kinds, the positive and the negative; opposite charges attract and like charges repel one another. From Coulomb I came into the person of Luigi Galvani, Professor of Anatomy and Gynecology at Bologna. In the course of my dissections, I observed that a dead frog could be made to twitch and convulse as if it were alive by any of three methods. First, it may be connected to a lightning rod, which gathers electricity from the air in stormy weather. Second, it may be connected to an electrical machine, that is, a machine consisting of a sphere of sulfur rubbed by a woolen cloth. Finally, it may be connected to two different metals and the metals brought into contact with one another. Such contact may initiate an electric spark, bringing with it apparent life.

That which is below corresponds to that which is above; just as two metals in contact may initiate an electric spark, so two mortals in verbal or literary contact may initiate a hermetic spark which passes from one to the other. I am that spark, passed from Galvani to Allesandro Volta, in whom I continued the Work. I found that a salt solution could substitute for the frog in Galvani's experiment. In the simplest possible terms, a voltaic cell consists simply of two different metals separated by a conductive solution. Voltaic cells may be connected in series to form a pile or battery. The more cells in the battery, the longer the spark produced; the longer the spark, the greater the electromotive force; and the unit of electromotive force is called the volt.

From Volta to Hans Christian Oersted, I discovered in 1820 that the current from a voltaic battery produces a magnetic field. Having witnessed a demonstration of Oersted's results at a meeting of the Acad魩 e des Sciences, the I-dea passed to Andr頁 mp貥, in whom I discovered a way to measure this current, the unit of which is now known as the ampere, or amp. Big batteries produced more current than small ones, but had the same voltage. I solved this puzzle in 1827 in the person of Georg Ohm when I announced the principle of resistance, which limits the current produced by a cell of a given voltage. The unit of resistance is now known as the ohm. With the concepts of EMF, current, and resistance clarified I was poised for my greatest advances in understanding the spark.

In no other single mortal did I make more headway in understanding electricity than in the person of Michael Faraday. Apprenticed to a book-binder at the age of 14, the I-dea of electricity was transmitted to him by an article in the Encyclopaedia Britannica, and from that moment I established a home in him. I found employment as the assistant

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to the great chemist, Sir Humphry Davy, who had made a brilliant start in the relationship of electricity to chemistry. In Davy and Faraday I came to understand that the positive and the negative are not two species of electricity. No, there is only one electrical entity; positive and negative simply represent a greater or lesser quantity of this One Thing. In 1831 I announced the principle of electromagnetic induction, the basis of the transformer, by which electrical current may be converted from one voltage to another. In 1834 I announced the laws of electrolysis, unifying the I-deas of electric charge and chemical amount. The conversion factor from electric charge to moles is now known as the Faraday constant.

My researches culminated in the first practical application of electricity in the person of Samuel Morse. In 1844 I demonstrated a device using the batteries of Volta and the electromagnetism of Oersted to transmit I-deas electrically from one place to another. This "telegraph" made it possible to communicate over vast distances almost instantaneously. Most of the practical applications of electricity for the next half-century would simply be variations on the telegraph theme: electrical current produced by batteries, controlled by switches, flowing through wires, and producing motion via electromagnets. And even after the advent of vacuum tubes, radio, and television these simple devices would not lose their spark.

A little more than a century after the invention of the telegraph I re-minded myself of the spark once more. A mortal child constructed his Halloween costume from cardboard boxes and tin foil; he wanted to be a robot. He hungered for unconventional playthings, winding copper wire around a nail to make an electromagnet, constructing a telegraph set from tin cans and scrap wood, dismantling radios and telephones to scavenge precious capacitors, resistors, and diodes. He grew impatient with the over-protectiveness of contemporary science books for children, devouring instead the boy-scientist[2] books of previous generations. He learned to tame the Model T ignition coil and the flyback transformer. And so I began my current incarnation.

In the person of Dunn I sit writing these words on a laptop computer. The device is powered by batteries; the keys on the keyboard are nothing more than switches; and the motors, relays, and pumps in my inkjet printer are all activated by electromagnets. Without these simple marvels, the transistors and integrated circuits would slumber in darkness. Truly there are few modern wonders, whether of art or artifice, which do not depend on the power of the spark.

Notes

[1] Reference [18], p. 239.[2] Reference [47] is an excellent book for boy- and girl-scientists capable of

overlooking the un-apologetic sexism of the era in which it was written.

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21.2. I did some growing up in the city of New Orleans. One of the most striking landmarks of the Big Easy is the Greater New Orleans Bridge. There is a toll to cross this bridge and we can go a long way toward understanding the chemistry of batteries by examining the economics of bridge tolls. Why would anyone pay a bridge toll? Obviously because they want to get to the other side. How much of a toll would they be willing to pay? It depends on how badly they want to cross the river by car; let us call this desire the automotive force, or AMF, with units of dollars/car. The AMF depends on the location of the bridge; it will be high when the bridge connects, for example, a major commercial district with a major residential area. People will pay a toll to cross the bridge when the cost of the toll does not exceed the automotive force.

Let us consider a chemical analogue. Whenever we have manipulated alkalis in this book we have been careful to avoid the use of aluminum containers or utensils. The reason is that aluminum is easily oxidized by alkaline solutions:

When sodium hydroxide is sprinkled on wet aluminum foil a vigorous reaction takes place; the solution foams up, the foil is eaten away, and a great deal of heat is released. If we were to look closely at the foil we would find some areas in which aluminum is being oxidized to aluminum hydroxide; these areas comprise the anode, the site of oxidation. In other areas water is being reduced to hydrogen gas; those areas comprise the cathode, the site of reduction. In the example so far, half of the aluminum foil acts as the anode and half as the cathode with electrons being passed from anodic regions to adjacent cathodic ones. Such a situation is analogous to a city in which residences and businesses are mixed in and among one another; there is plenty of traffic, but since it is mostly local it would be difficult to collect a toll from people commuting to and from work.

The situation would be different if we were to connect to the aluminum foil an inert metal, one which does not react with alkali. Such a metal could not be oxidized by the alkali and so it could not serve as an anode. But electrons would be able to move from the anodic regions of the aluminum foil, through the inert metal, and into the water. The water would be reduced to hydrogen gas at the surface of the inert metal and so the inert metal would serve very well as a cathode. With reduction taking place at the inert metal surface, the entire surface of the aluminum foil would become available for oxidation. Such a situation is analogous to a city in which the business district is adjacent to the residential district; there would be an automotive force reflecting the desire of people to travel from one district to another, but no way to collect a toll.

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Let us now separate the aluminum foil from the inert metal, leaving each one in contact with the alkaline solution. Such an arrangement, a voltaic cell, is analogous to a city in which the business and residential districts are separated by a river. Figure 21-1 shows a process schematic for the aluminum-alkali cell. Here the electrons consumed at the anode and produced at the cathode are shown explicitly as reactants and products. The electrons on the reactant side represent a wire dipping into the solution, or electrolyte; hydrogen gas is produced at this wire, the cathode. The electrons on the product side represent the wire leading from the aluminum anode. With the cathode at the top and the anode at the bottom of the schematic, we can imagine the electrons flowing downhill, impelled by the electromotive force, or EMF, analogous to the automotive force. Of course, in the physical cell the relative placement of the anode and cathode is irrelevant. The convention established in the process schematic simply helps us to re-meme-ber the direction of electron flow. If the wire from the anode is connected to the cathode we have, in effect, created a bridge and may "charge" the electrons a "toll" for the privilege of crossing over; we might require them to light a bulb or turn a motor. How much of a toll would they be "willing" to pay? Any amount up to the value of the electromotive force, measured in volts.

Figure 21-1. The Aluminum-Alkali Cell

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Equation 21-1. Four Electrochemical Reactions

The EMF, or voltage, of a cell depends on the identity of the conductors and on the identity and concentration of the electrolyte. Equation 21-1 shows the unbalanced reactions for four voltaic cells, (a) the aluminum-alkali cell, (b) the Leclanche carbon-zinc cell, (c) the alkaline carbon-zinc cell, and (d) the lead-acid cell. The aluminum-alkali cell is the one we have been discussing so far, with an EMF of 1.0 V, more or less, depending on the concentration of the alkali. The Leclanche cell is the familiar flashlight battery, with a voltage of 1.5 V. The alkaline cell comes in the same voltage and sizes as the Leclanche cell. Six lead-acid cells, at 2.0 V each, comprise the common automobile battery. Take a moment to balance these redox reactions using the method of Chapter 11 and then sketch out the process schematics for each cell. You can identify the anodes and cathodes by re-meme-bering that electrons go in at the top, down through the cell, and out the bottom; under this convention, the positive cathode is the electrode on top.

The flow of electrons is really only half the story of the voltaic cell. Returning to the bridge analogy, the Greater New Orleans Bridge collects tolls only from cars traveling towards the business district; the return trip is free. The "return trip" in our aluminum-alkali cell is afforded by the alkaline solution. Negative hydroxide ions are produced at the inert metal cathode and consumed at the aluminum anode. Just as the electrons flow from anode to cathode, the hydroxide ions migrate through the electrolyte solution from cathode to anode. This migration of charge through the electrolyte constitutes an electric current equal and opposite to that in the wire. The unit of electric current is the ampere or amp. The electric current is proportional to the rate at which electrons flow through the wire. The more electrons cross over in a given time period, the higher the current. Thus electric current is analogous to the rate at which traffic crosses the bridge.

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The amount of bridge traffic depends, of course, on the AMF; if people have little motivation to cross the bridge, traffic will be light. But it also depends on the population of the city; a large city will have more traffic than a small one. Similarly, the current that can be delivered by a voltaic cell depends on the size of the cell. For example, flashlight cells come in different sizes, D, C, A, AA, and AAA. Each of these cells has the same voltage. But the big D cell can deliver more current than the tiny AAA cell. Hence a boom box, with its large speakers and motors, requires D cells; a pocket radio, with only headphones, can get by with AAA cells.

The big payoff for the bridge comes when lots of people are willing to pay high tolls for long periods of time. For electrical devices, the payoff is called energy, with units of joules:

1 joule = (1 volt)(1 amp)(1 second)

The rate of energy delivery is called power, with units of watts:

1 watt = (1 volt)(1 amp)

Thus 1 joule = (1 watt)(1 second)

We can use these definitions to do UFA, for example:

Q: How many mega joules of energy are provided by the delivery of 1 kilowatt of power for 1 hour?

Q: How many mega joules of energy are provided by the delivery of 1 kilowatt of power for 1 hour?

A: 3.6 MJ

More power is delivered by two voltaic cells than by one and there are two ways of connecting cells to form a battery. When two cells are connected in series the anode of one is connected to the cathode of the other. For example, when Leclanche or alkaline cells are loaded into a flashlight the "nose" or button end, the positive cathode, of one is placed in contact with the flat bottom, the negative anode, of the next. The voltage of this battery of two cells will be double that of either cell on its own. You can think of two process schematics stacked one atop the other; the electrons "fall" twice the normal distance. For series batteries of more than two cells, the voltage is proportional to the number of cells. Thus, if you cut open a 9-volt transistor battery, you will find six 1.5-volt cells within. Similarly, a 12-volt automobile battery consists of six 2-volt lead-acid cells. Since the same current flows through each of the cells in a series battery, the power delivered is proportional to the number of cells; two cells deliver twice the power of one.

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Conversely, when two identical cells are connected in parallel, that is, anode to anode and cathode to cathode, the voltage of the pair is the same as that of an individual cell. Since the same current flows through each cell, the current of the pair is twice that of either on its own. Six 1.5-volt cells in parallel will have a voltage of 1.5 volts, but will be able to deliver more current than any of them alone. You can think of the process schematics arrayed side-by-side. In effect, the cathodes and anodes form one giant cathode and one giant anode. Since the voltage is the same as that of the individual cells but the current is proportional to the number of them, the power delivered is, again, proportional to the number of cells; two cells deliver twice the power of one. Thus more power is delivered by more cells no matter which connection scheme is used. The choice of series of parallel connection depends on the voltage required by the device to be powered.

This chapter has introduced a host of intangible concepts and their units: the volt, the amp, the joule, and the watt. It may have left you weary, feeling small. When tears are in your eyes I will dry them all. I am on your side. When times get rough and friends just can't be found, just remember that a voltaic cell is like a bridge over troubled water.

Material Safety

Locate MSDS's for aluminum (CAS 7429-90-5), soda ash (CAS 497-19-8), and charcoal (CAS 7440-44-0). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[1]

Your most likely exposure will be eye or skin contact. In case of eye contact flush them with water and call an ambulance. In case of skin contact wash the affected area with plenty of water.

You should wear safety glasses while working on this project. Leftover soda can be flushed down the drain with plenty of water. Used aluminum foil should be thrown in the trash. Used charcoal may be thrown in the trash or saved for reuse.

Research and Development

You should not remain ignorant if you are to proceed in the Work. Know the meanings of those words from this chapter worthy of

inclusion in the index or glossary. You should have mastered the Research and Development items of

Chapter 17 and Chapter 20. Know two new unit factors from this chapter and be able to use them

to convert among the units, volts, amps, and watts.

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Given the skeleton reactions for the aluminum-alkali, Leclanche, alkaline, and lead-acid cells, be able to produce balanced redox reactions and process schematics.

Know the hazardous properties of charcoal and washing soda. Know the years in which the voltaic pile and the telegraph were

invented.

Know the contributions of Franklin, Volta, Amp貥, Ohm, and Faraday to the understanding of the spark.

Notes

[1] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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21.3. For this demonstration you will require some aluminum foil to be used as an anode in your voltaic cell. You will also need a cathode. I am fortunate enough to have had mercury-silver cathodes installed in my mouth when I was a child. That way, wherever I am, I am able to construct my own batteries. Only one more thing is needed: an electrolyte. This I also keep in my mouth. If you are similarly equipped, roll up a ball of aluminum foil and pop it into your mouth. Move it from side to side, chewing it thoroughly. If the aluminum happens to connect with a cathode, you will perceive an electric shock. Congratulations! You have constructed a voltaic cell.

Such a cell is not entirely satisfactory, however. While it produces a decent voltage, the current produced is quite small. Similar proof-of-concept batteries may be constructed from coins and salt-water or lemon juice. It may even be possible to run a tiny electronic clock from one of these, as such clocks do not require much in the way of current. But for lighting light bulbs and running motors and closing relays we need a lot more current than can be provided by a battery made from coins. The key to high currents is electrode surface area.

The battery you will build is a variation on the one described in the previous section. I chose aluminum as the anode because it is commonly available and because its standard reduction potential is rather large and negative. Ordinary table salt may be used as an electrolyte, but I have found that washing soda, sodium carbonate, produces a cell with a higher current. In the previous section we discussed platinum as an inert electrode. Such an electrode allows water itself to serve as an oxidizing agent. Recalling that the oxidizing agent is, itself, reduced and that the site of reduction is called the cathode, we see that platinum served as the cathode. Platinum is expensive, however, and so we require a cheaper material for our cathode. For electrical applications, carbon often serves as "poor man's platinum." We would like a form of carbon with a large surface area so that electrons will be more likely to cross our wire bridge than to move directly from the aluminum to the electrolyte. An inexpensive, high-surface-area form of carbon is met with in activated charcoal.[1] Activated charcoal is porous, like a sponge, and is sold for use in water filters for aquariums. Ordinary charcoal is electrically conductive, but lacks the large surface area of activated charcoal. Since it is inert, it may be re-used to make generations of voltaic cells.

You will also need a convenient container for your battery and I find that a small stainless steel mixing bowls serve well. Three mixing bowls of the same size will allow you to construct a battery of two cells in series. The larger the mixing bowls, the greater the current from your battery, but even small, 2-cup bowls will power a small motor or telegraph sounder.

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Figure 21-2. The Aluminum-Alkali Battery

Begin by preparing a saturated solution of washing soda. Fill a mixing bowl half-full of water, add some washing soda, and stir until it is dissolved. Keep adding washing soda and stirring until no more will dissolve and solid sodium carbonate settles to the bottom of the bowl. Fill each of your remaining bowls about 1/4 full of activated charcoal. Add enough washing soda solution to each bowl to make a slurry with the activated charcoal. The charcoal will foam; air is displaced from the pores as the charcoal absorbs the solution. You now have two bowls, 1/4 full of charcoal and electrolyte. We are about to add the aluminum anode, but we need a way to electrically insulate the aluminum from the charcoal while still allowing the ions in the electrolyte to move about.

To do this, line each of the two bowls with two thicknesses of paper towel. This will keep the charcoal separate from the aluminum while allowing the electrolyte to flow between the two. Add enough washing soda solution to completely wet the paper towels. Now line each bowl with aluminum foil, taking care that the foil does not tear the paper towels and that it does not touch the mixing bowl beneath. You are now ready to connect your two cells in series.

Place one of your cells inside the other, just as if you were stacking the bowls for storage. The aluminum anode of the lower cell should make contact with the steel-charcoal cathode of the upper one. Empty the third mixing bowl of leftover washing soda and nest it within the upper cell. You now have three bowls nested one inside the other. If you were to look at their contents in cross section, it would go steel-charcoal-paper-

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aluminum-steel-charcoal-paper-aluminum-steel. Use both hands to press the upper and lower bowls together. By compressing the charcoal, you reduce its resistance and thereby increase the current produced by your cells. Finally, use a strong rubber band to hold the bowls tightly together.

Inexpensive volt-ohm meters are now widely available. Use a meter to measure the voltage between the lower bowl and the middle bowl, the middle bowl and the upper bowl, and lower bowl and the upper bowl. Record these voltages in your notebook and compare the voltages of the individual cells to that of the two of them in series. If your meter measures milli-amps as well, record the current output of each individual cell and compare them to that of the two in series. The voltage of your battery should be about 2 V and the current should exceed 50 mA.

There is a break-in period with these cells as aluminum hydroxide builds up around the anode. Be patient before concluding that your cells are not operating correctly. If one or both of your cells fail to produce any output, it is likely that the paper is torn, shorting anode to cathode; the electrons are able to move from anode to cathode without going through your meter. If both cells perform, but one performs better than the other, it is likely that the under-performing cell does not contain enough electrolyte. Dismantle the cell and add some more sodium carbonate solution.

The real test of your battery is whether or not it is strong enough to power an electrical device. You may try lighting a small flashlight bulb, running a small motor, or activating a telegraph sounder. You may even build your own electromagnet. Simply wind 50-100 turns of insulated wire around an iron nail. The wire must be insulated; "bell wire" or "magnet wire" can be found at hobby electronics stores. Magnet wire may also be salvaged from "dead" motors. With a knife, strip the insulation from the two ends of the wire; connect one end to the lower bowl and one to the upper bowl of your battery. When connected to your battery, your electromagnet should be strong enough to lift a paper clip.

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Figure 21-3. The Battery

Your battery will remain at full strength for hours but not for days. Eventually the aluminum anode will be consumed and the electrolyte will become saturated with aluminum hydroxide. The charcoal, however, is inert. When your battery is dead, or when you are tired of it, throw the aluminum foil and paper towels away. Drain the electrolyte from the charcoal and wash the solution down the sink. Wash the charcoal with a few portions of fresh water, drain, and allow the charcoal to dry. The mixing bowls and charcoal may be used to make more batteries later on. I would not, however, put the used charcoal in my aquarium filter; the residual alkali and aluminum hydroxide would probably not be good for fish.

The aluminum-alkali battery is far from optimal. One of the most practical cells of all time is the Leclanch 頣 ell, developed by Georges Leclanch 頩 n 1866. Batteries up to that time had consisted of jars filled with liquid, which were subject to breakage, spillage, and leakage. Leclanch 駳 innovation was to produce a "dry cell," which might more

accurately have been called a "moist cell." The Leclanch 頣 ell consists of a carbon/manganese dioxide cathode, surrounded by a paste of ammonium chloride, separated from the zinc anode by a layer of paper. The redox reaction is:

Zn(s) + 2 MnO2(s) + 2 NH4Cl(aq) = ZnCl2(aq) + Mn2O3(s) + 2 NH3(aq) + H2O(l)

Unlike the aluminum-alkali cell, the Leclanch 頣 ell produces no gases and so it may be hermetically sealed. As the cell ages, however, the ammonium chloride gets used up and the cell voltage drops. This drawback has been addressed by the alkaline dry cell.

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The alkaline dry cell uses the same anode and cathode as the Leclanch 頣 ell and consequently has very nearly the same EMF. Rather than using ammonium chloride as the electrolyte, however, it uses potassium hydroxide. The redox reaction is:

Zn(s) + 2 MnO2(s) + H2O(l) = Zn(OH)2(s) + Mn2O3(s)

Because the electrolyte is not consumed in the reaction, its concentration remains constant as the cell ages. Consequently the EMF of the alkaline dry cell remains nearly constant as it ages.

All of the cells we have discusses so far have been primary cells; once the anode is consumed in the reaction, the cell is dead and must be discarded. A secondary cell is one which can be recharged, that is, one in which the redox reaction may be run backwards to restore the original state of the cell. We will re-meme-ber this I-dea in Chapter 25.

Quality Assurance

Record in your notebook the voltages and currents produced by your individual voltaic cells and by the two cells connected in series. Your battery passes muster when it can operate an electrical device.

Notes

[1] In early versions of this project I used copper dishwashing pads as the cathode. The I-dea for using activated charcoal came to me from Reference [46].

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Chapter 22. Perkin (Aniline?Dyes)

The real inventor of practical gas-lighting is William Murdoch, who in 1792 lit his workshops at Redruth, Cornwall, with a gas obtained from coals. His operations remained unknown abroad for some ten years, and hence the French consider Lebon as the inventor of gas-lighting, since he lit (1801) his house and garden with gas obtained from wood. The first more extensive gas-work was established in 1802 by Murdoch, at the Soho Foundry, near Birmingham, the property of the celebrated Boulton and Watt; and in 1804 a spinning-mill at Manchester was lighted with gas. From that period gas-lighting became more and more generally adopted in factories and workshops, but not before the year 1812 did this mode of lighting become introduced into dwelling-houses and streets, a few of which in London were lit with gas in this year; while in Paris gas was first introduced in 1820. From that year gas-lighting may be said to have become of general importance in Europe, and now there is hardly any important place on the Continent where it is not in use, while as regards the United Kingdom in no portion is gas-making and lighting so general over town and country as in Scotland.

— Wagner's Chemical Technology, 1872 AD [1]

22.1. Gas! Now that's a business to be in.

I don't believe it! Here I am, supposedly the representative of the alchemical element air, waiting patiently for a chapter that might possibly have some remote connection to the element for which I am supposed to speak, and when one finally comes around the Author assigns it to a confabulating bug with about as much air-worthiness as a lead Zeppelin. And it makes me wonder.

If there's a bustle in your hedgerow, don't be alarmed now. This chapter isn't about gas; it's really more about tar. See, folks had been making charcoal from wood since God was a child. Now, to make charcoal, you put wood in a sealed container, and heat the bejeezus out of it in the absence of air. Bejeezical water goes up the smoke stack, like it shows in Equation 1-1, and the stuff that's left behind is charcoal, which is mostly carbon. That charcoal burns hotter and cleaner than wood and so it's great for smelting metals and all.

Now all this wood burning caused a scarcity of wood in Europe, a kind of energy crisis. Coal was an obvious alternative since it, like charcoal, consists primarily of carbon. But coal wasn't any good for smelting metals because it contains impurities like sulfur which weaken metals. It wasn't until 1603 that a fellow named Hugh Platt tried heating coal anaerobically, driving off bejeezical gases and producing purified carbon in

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the form of coke. Coke was a big hit with steel-makers, but it would be another two hundred years before anyone got the bright idea to do something with the gas from the coke ovens.

That's where Phillipe Lebon comes in. See, he grew up in a charcoal-producing region of France and was inspired to use the gas from the charcoal ovens for illumination. He got a patent in 1799 for his "Thermolamp" and even thought up an engine to run off of gas; he was way ahead of his time. Who knows what he would have come up with if he hadn't been stabbed to death in 1804?

Something quite remarkable, no doubt, but returning to events that actually materialized, William Murdock began experimenting with gas from the coke ovens, lighting his workshops in 1792. When news of Lebon's work reached Murdock's employer, Matthew Boulton, they pushed forward with commercial development of gas lighting. Murdock described the basic principles of gas lighting to the Royal Society in 1808.

The recovery of gas from a coke oven may be viewed as a kind of distillation. In the distillation of Chapter 16 a mixture is simply separated into its components, for example, a solution of alcohol and water is separated into alcohol and water. Turn up the temperature, however, and the material you are distilling may start to decompose. Wood, for example, may decompose into charcoal and water. In such a destructive distillation the materials recovered were not present at the beginning, but were produced during the distillation. The destructive distillation of coal produces four products which are easily separated from one another; coke remains in the pot; gas driven from the pot is passed through a condenser where some of it condenses into an aqueous solution of ammonium carbonate; separating from this solution is a tar which is insoluble in water; the remaining gas is ready for distribution as illuminating gas.

But like I said at the beginning, this chapter is not so much about the gas; it's about the tar. That tar is a complicated mixture of compounds and folks started to wonder what they were and what could be done with them. In 1825 Michael Faraday isolated benzene from the destructive distillation of whale oil. Pretty soon everybody and her dog were destructively distilling things to see what would come of it. In 1826 Otto Unverdorben isolated aniline from indigo and in 1834 Friedlieb Runge isolated aniline and phenol from coal tar. So with a second, non-destructive distillation folks started producing chemicals from coal tar, chemicals like benzene, toluene, aniline, phenol, and naphthalene. Folks figured that since these compounds came from the destruction of organic compounds, maybe it would be possible to make good and useful things by putting them back together again.

Gas had become a major industry—

—but this chapter is more about tar. See, at that time malaria was a problem for the British Empire, on account of its rampant colonialism, and the best treatment for malaria was quinine, extracted from the bark of a South American tree. To promote chemical

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innovation the British started up the Royal College of Chemistry in 1845, with the great German chemist, Wilhelm Hofmann, as director. In 1856 Hofmann's assistant, William Perkin, set out to synthesize quinine from coal tar. Perkin reacted aniline with potassium dichromate, a really strong oxidizing agent; the resulting black goo was definitely not quinine, but it made a beautiful purple solution in alcohol. Perkin called it "mauveine" and dropped out of college at the age of 18 to develop his new synthetic dye. British dyers didn't think that mauve would catch on, but the Paris fashion houses liked it so much that Perkin was able to retire at the age of 36. I told you it was a great business.[2]

And—

—once folks knew that it was possible to make artificial dyes from coal tar, new dyes started coming out of the gasworks, so to speak. Variations on the original synthesis produced dozens of dyes from aniline: aniline reds, aniline violets, aniline greens, yellows, browns, and blues. Substituting phenol or naphthalene for aniline produced two more distinct families of artificial colors. There seemed to be few colors which could not be fashioned by art and ingenuity from coal tar.

Hofmann's leadership at the Royal College had given the British a head start in synthetic dyes, but when Hofmann returned to Germany in 1865 British dyestuffs dropped the ball. In France, synthetic dye manufacturers charged so much for their products that demand there shifted to natural and imported dyes. But in Germany, the situation was ideal for the development of these new wonder dyes. State-subsidized technical schools fed talented students to the universities; the universities promoted collaborative work, sending students to study in Britain and France; and universities cooperated with industry on joint research projects. Furthermore, with thirty-nine states in the German Confederation, patent enforcement was a problem and the resulting competition produced lean, mean, dye-making machines. German unification in 1871 made patent enforcement easier, but by that time the big players in German dyestuffs were off and running; Badische Anilin und Soda Fabrik (BASF) was out of the gate in 1861, followed by Farbwerke Hoechst in 1862; Freidrich Bayer rounded the bend 1863, followed by Kalle & Co in 1864; and a late starter from 1867 was reorganized in 1873 as Aktien Gesselschaft f?lin Fabrikation (AGFA). Throughout a century of wars, depressions, revolutions, and mergers Bayer, Hoechst, and BASF were the top dogs of the chemical industry.

I guess you could say, "They were buy-y-ing a ta-ar-way to heaven."

Notes

[1] Reference [29], pp. 645-646.[2] For a popular account of Perkin's career, see Reference [59]. For the early history of

the synthetic dye industry, see Reference [74].

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22.2. Tar happens. It happens in a fireplace; as the wood is heated it begins to break down into charcoal and a kind of greasy steam which condenses in the chimney. That greasy steam is tar. Tar happens in a cigarette; as the tobacco is heated, bejeesical nicotine passes through the filter. The stuff trapped by the filter is tar. Tar happens when the fat from your barbecued chicken drips on the hot coals; the fat decomposes and tar is deposited on the inside lid of the barbecue grill. Tar happens whenever things burn or scorch, releasing vapors which later condense somewhere else. If they condense in your nose, you smell them. Tars make your clothes smell of cigarettes after you leave the bar. Tars alert you that the bacon is burning. Tars make up the aroma of chestnuts roasting on an open fire. Tar is not a single substance; it's just a fancy catch-all name for all the various substances produced by heating things until they decompose. The fancy name for heating things until they decompose is destructive distillation.

Destructive distillation was all the rage in the 1820's. In 1826 Otto Unverdorben destructively distilled indigotin, the blue dye we met in Chapter 12. You probably just got used to molecular formulas like C16H10N2O2 and empirical formulas like C8H5NO when Chapter 16 went and introduced the molecular model as a way of showing how the atoms in a molecule are hooked together in 3D. Figure 22-1 shows a molecular model for indigotin. Molecular models are great when you have a wooden or plastic model or a computer program to draw them, but when you need a quick 2D sketch of how the atoms are connected, chemists use a structural formula, like the one shown at the bottom of Figure 22-1. To understand structural formulas you need to realize that carbon atoms generally make four bonds. So whenever four lines come together in a structural formula, that point stands for a carbon atom, even if no "C" is shown. A curious feature of indigotin is that it contains two six-member carbon rings with alternating single and double bonds. Those rings are unusually strong on account of the double bonds and so when indigotin decomposes, those rings are likely to survive in the decomposition products. You can imagine lots of ways for the indigotin molecule to fall apart, but the bonds cut by the dotted lines in Figure 22-1 are among the most likely to break during destructive distillation. When that happens, the two hunks containing the carbon rings turn into two molecules of aniline, the stuff Unverdorben recovered in 1826.

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Figure 22-1. Indigotin

What a surprise it must have been when Friedleib Runge isolated this same aniline from the destructive distillation of coal. Granted, most of the coal turns into coke, just as wood turns mostly into charcoal. In addition, several gases are produced—hydrogen, methane, and carbon monoxide. Water and ammonia are also produced. But the distinctive aroma of burning coal comes from the mixture of compounds which constitute coal tar, those volatile components which become gases at elevated temperatures but which condense into liquids and solids as they cool back to room temperature. The strong smell of these coal-tar products has given a name to a whole class of organic compounds: the aromatic compounds. One hundred kg of coal will yield approximately 3 kg of tar; distillation of that tar will produce about 1 kg of aromatic compounds.

You might think that aromatic means smelly, and surely that was the original meaning. But as time when on, it turned out that the aromatic coal tar distillates all shared a common structural feature, that curious six-membered carbon ring. As more and more compounds containing this ring were identified, it turned out that not all of them were smelly. But the name aromatic stuck fast like a spider. The meaning of the term is a little more technical these days, but for our purposes, an aromatic compound is one which contains the six-membered, alternately single and double-bonded carbon ring, the aromatic ring.

Each carbon atom of an aromatic ring can be bonded to one other atom. When all six carbon atoms are bonded to hydrogen atoms, we have a molecule of benzene, shown in Figure 22-2(L). When I was in college benzene was a common laboratory chemical, dry-cleaning solvent, and paint thinner. Benzene is not particularly toxic by ingestion, with an LD50 of 3.8 g/kg in the rat. Because of its low boiling point, inhalation is more likely to be a problem than ingestion, but with adequate ventilation acute exposure to benzene is

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no more hazardous than exposure to many solvents currently available at hardware stores. But in addition to the risk of acute exposure, there is a danger of chronic exposure; occupational exposure to benzene is associated with an increased risk of cancer. Consequently, benzene has been replaced by toluene for household and many commercial uses.

Figure 22-2. Benzene, Toluene, Aniline

Toluene, shown in the middle of Figure 22-2, is familiar to most people as the solvent in model airplane glue. You might think of it as a benzene molecule with an extra carbon atom attached. As with benzene, inhalation is the most important hazard but with adequate ventilation it can be handled safely. Unlike benzene, chronic exposure to toluene is not associated with cancer and toluene can still be found in hardware stores among the paint thinners.

Aniline, shown on the right of Figure 22-2, is the stuff originally produced by the destructive distillation of indigotin. You might think of it as a benzene molecule with an ammonia molecule attached. Aniline is an oily brown liquid only slightly soluble in water. Though aniline is the compound needed for the synthesis of mauveine, indigo is too expensive to be used as a commercial source and coal tar doesn't contain enough of it to be useful. Since benzene is the most abundant coal-tar aromatic it is natural to look for a way of making aniline from benzene. This synthesis occurs in two steps; first benzene is oxidized by nitric acid:

C6H6 + HNO3 = C6H5NO2 + H2O

Second, the nitrobenzene produced in the first reaction is reduced by iron filings:

C6H5NO2 + H2O + 2 Fe = C6H5NH2 + Fe2O3

Similar reactions produce toluidine, C7H8NH2, from toluene. There is a hitch, however, in that the NH2 might appear in any of three positions on the ring. Consequently, there are three "flavors," or isomers of toluidine: ortho-, meta-, and para-toluidine, shown in Figure 22-3. Each of these isomers has its own boiling point and other physical properties. Oxidation of toluene with nitric acid and subsequent reduction with iron yield

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predominantly ortho-toluidine and para-toluidine. Early coal-tar distillation did not completely separate benzene from toluene and consequently commercial "aniline" contained significant quantities of the toluidines. Just as organic molecules which contain the OH group are classified as alcohols, those containing the NH2 group are classified as amines. The aromatic amines, aniline and the toluidines, are the raw materials for the synthesis of the dye, mauveine.

Figure 22-3. The Toluidines

When you oxidize pure aniline in an acidic solution, the product is the dye pseudomauveine:

4 C6H5NH2 = C24H19N4+ + 9 H+ + 10 e-

The remnants of the four original aniline molecules are evident in the structure of pseudomauveine, as shown in Figure 22-4. The "R's" in the structure are simply hydrogen atoms, for the moment. Note that pseudomauveine is a cation and so there must be an anion in the solution to balance it. If the reaction took place in acetic acid, the result is pseudomauveine acetate; if it took place in hydrochloric acid, the result is pseudomauveine chloride, etc. No matter which acid was used to make it, pseudomauveine is a brown dye, not very soluble in water but quite soluble in ethanol. Brown dyes are neither very rare nor much sought after and if Perkin had used pure aniline in his attempted synthesis of quinine, he would not have been much impressed with the results.

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Figure 22-4. Pseudomauveine and Mauveine

Perkin's aniline was not pure, however; it contained o-toluidine and p-toluidine. With these other amines in the mix, the "R's" in Figure 22-4 might turn out to be either hydrogen atoms or CH3 groups, depending on the mix of aniline and the toluidines. With CH3 groups at Ra and Rc, two slightly different purple dyes result, one (mauveine A) with Rb = H and the other (mauveine B) with Rb = CH3. Like pseudomauveine, the mauveines are not very soluble in water but are quite soluble in ethanol. Unlike pseudomauveine, the mauveines are purple, a color rare among natural dyes and prized as a status symbol. If you had found a way to make such a color in a world of browns and blues and yellows and tans, you would very likely have dropped out of school as Perkin did. And that color would have made you as wealthy as the folks already wearing purple clothes.

Material Safety

Locate MSDS's for ethanol (CAS 64-17-5), aniline (CAS 62-53-3), o-toluidine (CAS 95-53-4), p-toluidine (CAS 106-49-0), acetic acid (CAS 64-19-7), and sodium hypochlorite (CAS 7681-52-9). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[1]

In the course of this project you will combine vinegar, aniline, and laundry

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bleach to produce a solution of dye in ethanol. By themselves, vinegar and bleach would release chlorine gas, but you will be using only a few drops of bleach and the liberated chlorine will immediately oxidize the aniline; you are more likely to smell chlorine from the bleach itself than from your dye solution. Be aware that sodium hypochlorite will bleach clothing.

You should wear safety glasses and rubber gloves while working on this project. All activities should be performed in a fume hood or with adequate ventilation. Leftover aniline and toluidine should be converted to dye as described below. Leftover dye should be washed down the drain with plenty of water unless prohibited by law.

Research and Development

Before you get started, you should know this stuff. You better know all the words that are important enough to be

indexified and glossarated. You should know all of the Research and Development stuff from

Chapter 18 and Chapter 19. You should know the structures and properties of the coal-tar

distillates: benzene, toluene, and aniline, as shown in Figure 22-2. You should know the structures of aniline and o-, m-, and p-toluidine

as shown in Figure 22-3. You should be able to recognize the remnants of aniline and toluidine

in mauveine and pseudomauveine, as shown in Figure 22-4. You should know the hazardous properties of ethanol, acetic acid,

aniline, o-toluidine, p-toluidine, and sodium hypochlorite.

You should know who started the synthetic dye industry, when he did so, and which modern companies were founded as dye companies.

Notes

[1] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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22.3. It has been fun researching this book and trying to figure out how to make things from scratch. Sometimes I have gained inspiration from modern laboratory manuals, but more often than not, the projects have developed by trying to imagine what commonly-available materials could be used to make something appropriate to each chapter; this chapter was one of the last ones to "come together." At first I imagined that I would make mauve as Perkin had done, but the more I learned about Perkin's synthesis, the more discouraged I became. Perkin dissolved his aniline in aqueous sulfuric acid. Oxidation of this solution with potassium dichromate produced an insoluble brown goo which would have to be filtered, and filtering goo is no picnic. Once filtered, the goo would be extracted with ethanol, which would dissolve the dye. Perkin's yield was on the order of 5%, that is, 5% dye, 95% useless goo. Could caveman chemists be expected to make such a dye when even experts got such poor yields? Furthermore, chromium compounds pose a health risk and a disposal problem I was not sure I wanted to impose on beginners. Finally, neither aniline, nor toluidine, nor potassium dichromate is a household item. It seemed that there was much to be said against mauve as a project.

Nevertheless, I wondered whether it might be possible to substitute a household oxidizing agent for potassium dichromate. Wagner's Chemical Technology of 1872[1] reported that bleaching powder, calcium hypochlorite, had been used as an alternative to potassium dichromate. I wondered whether sodium hypochlorite, modern laundry bleach, might also do the trick.

I also wondered whether the goo-filtering step might be avoided by performing the oxidation in ethanol rather than aqueous solution. I expected that the goo might settle to the bottom of the container and the ethanolic dye might simply be poured off. My very first try using laundry bleach in ethanol took my breath away; it resulted in a rich purple dye and no goo at all. Not only did the household chemical work, it apparently worked better than any combination I had found in the chemical literature! But it was when I applied the dye to some silk cloth that I knew how Perkin must have felt.

To begin with, you need some aniline. If you buy aniline from a chemical supply it will likely be of higher purity than you need. Pure aniline will produce the brown pseudomauveine which, while it is a perfectly good brown dye, is unlikely to have the visual impact of mauveine. To make mauveine you need aniline contaminated with o-toluidine and p-toluidine. The optimum proportions are one mole aniline, two moles o-toluidine, and one mole p-toluidine. Since these three compounds have similar molecular weights and densities, you can make it up as 1 gram of aniline, two of o-toluidine, and one of p-toluidine. I will refer to this mixture as "aniline oil." The four grams just described will make a lot of dye; in a classroom situation, this amount of aniline oil will be enough for about a hundred students.

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You will make two solutions. First measure one mL of white household vinegar using a graduated pipette or a graduated cylinder. Place the vinegar into a small vial and add 1 drop of aniline oil to it. The oily aniline will dissolve completely within a minute or two. Now measure 5 mL of ethanol into a graduated cylinder and add 10 drops of bleach to it. Add the ethanol-bleach solution to the vinegar-aniline solution and put the cap on the vial. The combined solution should immediately turn brown as the aromatic amines are oxidized. The brown color will turn blue in a few minutes and then purple in an hour or so. This vial of dye is sufficient for dyeing a 12-inch silk handkerchief.

Is your dye a single compound or a mixture of compounds? To answer this question we'll use a technique invented by Otto Unverdorben for precisely this purpose: paper chromatography. You've probably seen water-soluble ink bleed when a piece of paper gets wet. If the ink is a mixture and if the mixture's components have different solubilities, then the more soluble components bleed faster than the less soluble ones. As the water spreads out on the paper, the more soluble ink components separate from the less soluble ones. This is the main idea behind chromatography. Since our dye is not soluble in water we'll use a solution of vinegar and ethanol as a solvent.

To make your chromatograms you'll need a piece of filter paper (90 mm in diameter), a pencil, a glass capillary tube (1 mm inside diameter, open both ends), a glass jar with a lid, and a solvent consisting of 75% white vinegar and 25% ethanol. Begin by folding the filter paper lengthwise into quarters and snipping off the ends with a pair of scissors. The resulting paper will be able to stand up, as shown in Figure 22-5(L). With a pencil[2] make a dot 1 cm from the bottom of the paper. Your goal is to place as much dye onto the smallest spot possible. To this end, dip the capillary tube into your dye; dye will climb into the tube by capillary action. Practice using the capillary tube by touching it briefly to a piece of paper towel. Some of the dye will be transferred to the towel, forming a spot. Practice making the smallest possible spots. Once you have some confidence, draw up some more dye into the capillary tube and spot the paper towel until a 1 cm column of dye remains in the tube. You will now deliver that 1 cm column of dye onto the pencil dot on your filter paper, as shown in Figure 22-5(R). Place the first spot onto the pencil mark and count to thirty, allowing the spot to dry. Place a second spot on top of the first and count to thirty. Continue to build this spot until you have delivered the entire 1 cm column of dye onto the filter paper. If you do a good job this "young" dye spot will be no more than 5 mm in diameter. If you mess up, just spot another paper until you get it right. You will now develop your chromatogram using a solvent composed of vinegar and ethanol.

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Figure 22-5. Folding and Spotting the Chromatogram

Mix 9 mL of white vinegar and 3 mL of ethanol in a glass jar large enough to contain your filter paper. Screw the lid onto the jar and shake it to thoroughly wet the walls of the jar. We want the air in the jar to be saturated with solvent so that the solvent doesn't evaporate as it climbs the paper. Open the jar and place your filter paper inside, standing it on the end with the spot and taking care that the paper doesn't touch the walls of the jar. Then replace the lid. The vinegar/ethanol solvent will climb the filter paper, taking the dye with it. Your chromatogram will take about an hour to develop, as shown in Figure 22-6(L).

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Figure 22-6. Developing and Interpreting the Chromatogram

When the solvent has climbed to within 5 mm of the top, remove your filter paper from the jar and draw a pencil line where the solvent stops. Your chromatogram will look something like Figure 22-6(R), which is shown in color on the back cover of this book . The chromatogram on the right is for a dye which has been allowed to mature for a week. At least four substances are evident in these chromatograms. Spot a was present in the young dye but not in the mature one. It represents a precursor, a combination of two or perhaps three aromatic amines. Spots b and c overlap but represent at least two different mauve dyes with slightly different colors. Spot d is soluble in the highly ethanolic dye solvent, but not in the less ethanolic chromatographic solvent. Consequently, it has not budged from the original spot.

Chemists characterize chromatographic spots by comparing the distance traveled by each one to that traveled by the solvent. Measure the distance from the original pencil dot to the center of each spot and to the solvent line. Divide the distance for each spot by that for the solvent to get what is called the Rf value (ratio of fronts). Spot a in the chromatogram shown has an Rf value of 56/65 = 0.86. Spot d has an Rf of 0.

By the time your chromatogram develops, your dye is mature enough to dye your silk handkerchief.[3] Push the entire handkerchief into your vial of dye and once it has taken up the dye, fish it out again. Spread it out on a few sheets of newspaper and allow it to dry. Be careful not to drip dye where it ought not to be. Once it's completely dry, you can wash your handkerchief with soap and water. Wow, is that a pretty color!

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Quality Assurance

Your dye is not so great if the color isn't fast. If the color hasn't washed out, tape your mauve handkerchief and your chromatogram into your notebook as proof of your dye-making prowess. Make a table of Rf

values for your dye spots and explain how many substances were present in your young dye.

In this way was the World created. From this there will be amazing applications because this is the pattern.— The Emerald Tablet of Hermes Trismegistos

Notes

[1] Reference [29], p. 577.[2] A mark made with a pen would bleed.[3] If you wish, you can hold back some of your dye and repeat the chromatography

after it has matured for a week.

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Chapter 23. Eastman (Photography)

It has already been observed, that when a solar spectrum falls upon paper covered over with chloride of silver, the chloride turns black in the more refrangible regions, and from this and similar experiments we have been lead to the knowledge that there exists in the sunbeam a principle which can bring about chemical changes.

This fact has been perceived from the beginning of the present century; but, of late, much attention has been given to these rays, and from a consideration of the phenomena they exhibit, I have endeavored to prove that they constitute a fourth imponderable principle of the same rank as heat, light, and electricity; and for the purpose of giving precision to this view, have proposed that they be called Tithonic rays, from the circumstance that they are always associated with light; drawing the allusion from the classical fable of Tithonus and Aurora. …

The process of the Daguerrotype depends on the action of the Tithonic rays. It is conducted as follows: A piece of silver plate is brought to a high polish by rubbing it with powders, such as Tripoli and rotten-stone, every care being taken that the surface shall be absolutely pure and clean, a condition obtained in various ways by different artists, as by the aid of alcohol, dilute nitric acid, &c. This plate is next exposed in a box to the vapor which rises from iodine at common temperatures, until it has acquired a golden yellow tarnish; it is next exposed, in the camera obscura, to the images of the objects it is designed to copy, for a suitable space of time. On being removed from the instrument, nothing is visible upon it; but on exposure to the fumes of mercury, the images slowly evolve themselves.

To prevent any further change, the tarnished aspect of the plate is removed by washing the plate in a solution of hyposulfite of soda, and finishing the washing with water; it can then be kept for any length of time. …

There are many other photogenetic processes now known; several have been invented by Mr. Talbot; among them may be mentioned the calotype. Sir J. Herschel, also, has discovered very beautiful ones, and these possess the great advantage over Daguerre's that they yield pictures upon paper. In minuteness of effect they can not, however, be compared to the Daguerrotype.

— John Draper, Textbook on Chemistry, 1850 AD[1]

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23.1. Well, where is the Figment? This is his chapter, is it not?

Looks like he's out of the picture.

When I look back on all the crap he's had from you two, it's a wonder he shows up at all. I guess your lack of admiration must have hurt him some; he could see the writing on the wall.

Speaking of writing on the wall, I guess that photography really started with the invention of the camera obscura, which is basically just a dark little room with a hole at one end. Light from outside comes in through that hole and projects an image of the outside scenery on the wall. Folks had been using the camera obscura since the sixteenth century, tracing the projected image onto paper with a pencil.[2]

All well and good, but you cannot consider a pencil drawing as a photograph. More fundamental to the birth of photography was my discovery in 1725, in the person of Johann Schulze, that silver nitrate darkens upon exposure to light. In modern photography it is from exposed silver salts that the photographic image is ultimately formed.

Formed, yes, as long as this image remains in the dark of the camera. But bring it out into the daylight and the whole shebang turns black. You are probably wondering how this image may be "fixed" permanently. I will tell you. In 1819 John Herschel discovered that aqueous sodium hyposulfite dissolves silver salts, most of which are otherwise insoluble in water. So by soaking your image in such a solution, the unexposed silver salts are washed away while the exposed areas remain. Voila! A permanent photographic image.

Now hold on there, cowboy. That's a heck of a theory, but it's still not a photo. Seems to me that the prize for first permanent photo goes to Nic 鰨 ore Ni 鰣 e for his 1826 snapshot of the view outside his window.

Snapshot, indeed! The exposure took eight hours.

So? Anyhow, in 1829 Ni 鰣 e hooked up with Louis Jacques Daguerre, who was wowing theater audiences with his elaborate special effects. Daguerre's "Diorama" involved projecting light through huge semi-transparent paintings, the motion of the lights and the paintings set to music. Daguerre used the camera obscura to sketch his huge paintings and looked to Ni鰣 e for inspiration on eliminating the tedious sketching.

Ni 鰣 e died in 1833, but not before Daguerre had caught the "photo bug." In 1837 he

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succeeded in producing a photo on silver-plated copper, developing the image with mercury vapor and fixing it with a solution of ordinary salt. All with an exposure under 20 minutes.

Impressive to him at the time, no doubt, but still not a photograph as we understand the term today. Both Ni鰣 e and Daguerre produced images on opaque metal plates. With such a system there is no way to copy the image, no way to enlarge it, no way to project it onto a screen. Without duplication and projection the I-dea of the motion picture would not have materialized; without motion pictures, no television; without television, no computer monitors. The metal-plate image is simply a fluke, an evolutionary dead end. I produced the first true photograph in 1835 in the person of Fox Talbot. This 10-minute exposure on translucent paper of the view through the windows of Lacock Abbey resulted in an image which could be projected, enlarged, and reproduced; all fundamental properties of photographs as we know them today.

An image which, like Daguerre's of 1837, deteriorated over time because silver salts are generally insoluble in water. In both instances the residual silver salts darkened with time, spoiling the original image. What can be a more important property of a photograph than that it should be permanent? I must remind you that Herschel discovered the solubility of silver salts in sodium hyposulfite in 1819 and I must further point out that he informed both Talbot and Daguerre of this in 1839, marking the beginning of true photography.

A step forward, to be sure.

It was a hum-dinger. Thanks. But to get back to the story, these Talbot prints were pretty fuzzy, as opposed to daguerreotypes, which were sharp as a tack. And what's more, talbotypes were funky, in that the parts of the photo which ought to be black were white and the parts that ought to be white were black. When you think about it, the talbotype image had everything except what was supposed to be there. In the daguerreotype, the parts of the photo which ought to be white were shiny silver and the parts that ought to be black were dark copper. Awesome! Napoleon III liked the daguerreotype so much that he gave Daguerre a nice pension and released the process to the public as a benefit to the human race.

Except in England; you were careful to patent the daguerreotype in Talbot's homeland. And for your information, the "funkiness" you complain about is called a negative image. It is a feature, not a problem. When you expose a second paper through the negative, the resulting print is a positive. My further research revealed that the image forms even before it is visible and that this latent image can be developed using gallic acid. I patented as this calotype process in 1842 and was able to use exposure times of 30 seconds, making portraiture possible.

That is, if you like fuzzy, funkified portraits. Plus, you had to pay for a license to produce calotypes, whereas daguerreotypes were patent-free (except in England). It just

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sounds too complicated to me, what with invisible negative latent images and all. Most folks agreed, and the daguerreotype business really took off, leaving the calotype in the dust.

Of course, the fuzziness of the calotype stemmed from the fact that the image penetrated the paper. A sharper negative image would result if it were produced on a glass plate rather than on paper, a suggestion Herschel made in 1839.

I was getting around to that; as Frederick Archer I formed an emulsion of silver salts and collodion on a glass plate in 1850. This thin emulsion enabled the production of sharp negatives. And in that same year, in the person of Louis Blanquart-Evrard I applied a similar technique to paper prints; I mixed silver salts with egg whites, forming a thin emulsion which coated only the surface of the paper. This "albumin paper," in combination with the collodion-glass negative, produced prints as sharp as the daguerreotype. With no way to duplicate or project daguerreotypes, it was only a matter of time before they went the way of the gaslight.

Well, it was a good business while it lasted. Eventually everybody and his dog had a somethingtype process named after him, what with little improvements here and there. The daguerreotype had made photography practical enough that pretty much anybody who wanted a picture could have one taken. But even the new-fangled somethingtypes had to be developed right after exposure, so unless you had a photographic darkroom, you couldn't take your own pictures. Then in 1889 a fellow named George Eastman, in the practical down-to-earth spirit of Daguerre, introduced celluloid roll film and a simple camera, the Kodak. For the first time you could take a whole roll of pictures and then send your film off to be developed and printed. From then on, it was "Momma don't take my Kodachrome away."

Notes

[1] Reference [11], pp. 90-93.[2] For a general introduction to the history of photography, see Reference [70].

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23.2. The photographic process is fundamentally a redox process. While the redox chemistries of iron and chromium have been explored from time to time, most photography, for most of its history, has been based on silver. It is most likely that every photograph you have ever seen has been a silver photograph. We saw in Chapter 9 that silver is primarily a by-product of lead smelting, in whose ore, galena, it is a common contaminant. The production of photo-chemicals begins with the oxidation of metallic silver with nitric acid:

Silver nitrate is, of course, soluble in water. To make a photographic emulsion we need an insoluble salt, one which will not leech out of the emulsion into the paper. You will recall from Chapter 7 that silver chloride is insoluble in water:

Silver bromide and silver iodide are also insoluble in water and these salts are more light-sensitive than the chloride. Consequently they are the silver salts of choice for modern photographic emulsions. But because sodium chloride, ordinary table salt, is easier for the amateur to obtain, we shall use it in the production of our emulsion.

If our paper were simply saturated with silver chloride, the resulting image would penetrate the paper resulting in a fuzzy image. To demonstrate this, write on a paper towel with a pencil; notice that the pencil line is sharp and that it does not penetrate or bleed into the paper. By contrast, write on a paper towel with an ink pen; notice that the ink soaks into the paper resulting in a fuzzy line. Writing paper is sized, that is, it is treated with substances to prevent ink from penetrating the paper. We shall use a similar trick to ensure that our silver chloride resides only at the surface of the paper.

To complete the photographic emulsion we need something which we can apply to the paper but which will be impervious to water once it is dry. This substance should bind well to the paper and remain flexible after processing. Collodion and gelatin have been used for this purpose, but in the spirit of using as many household substances as possible, let us choose egg whites. Egg whites contain the protein albumin, which is fairly impervious to water once it is dry, an observation which will be confirmed by anyone who has tried to wash dried eggs from the side of a house. Briefly, a mixture of sodium chloride in egg whites will be applied to paper and allowed to dry. Once dry, the coated paper will be sensitized with silver nitrate solution. Silver nitrate will react with sodium chloride to produce insoluble silver chloride which will reside only in the emulsion because the dried albumin will not allow it to soak into the paper. The sensitized paper must be stored in the dark until it is used.

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Now, silver ion is an oxidizing agent, that is, it is easily reduced to metallic, elemental silver. Albumin is a mild reducing agent, that is, it is easily oxidized. Place an oxidizing agent with a reducing agent and you are just asking for a redox reaction to happen. Given a day or two, the silver ion will be reduced to metallic silver:

As the silver is reduced the albumin will be oxidized to, well, to whatever it is oxidized to. Remember, proteins are polymers of amino acids; they form large, complex molecules and because the oxidation of the albumin has little to do with the resulting image, let us leave it at that and concentrate on the silver.

In the sensitized paper silver ion will be slowly reduced to metallic silver. You might expect, then, that after a few days the sensitized paper would look like a paper mirror. But you probably know from experience that powdered metals generally look black; take a file to any piece of metal and the filings will be dark gray or black. Silver is no exception. The reduced silver in the paper is not one big piece of silver to be polished to a mirror finish; it is in the form of tiny grains of silver. As a consequence, a sensitized paper left to its own devices for a few days will turn dark brown, gray, or even black.

Now for the magic! When light shines on the silver chloride/albumin emulsion the redox reaction happens more quickly than it does in the dark. Imagine now a sensitized sheet of paper, half of which is exposed to bright sunlight and the other half of which is covered up with an opaque card. The silver ions in the exposed area will be reduced to black metallic silver in a matter of minutes; the silver ion under the card will remain colorless or white. Imagine now that you remove the opaque card; what will you see? The half of the paper that was in the light will be black and the half that was in the dark will be white. This reversal of light and dark is referred to as a negative image.

But as soon as you remove the opaque card from the sensitized paper, the formerly unexposed white half will begin to turn black. In order to fix the image, we need to remove the light-sensitive silver chloride. Just washing it in water won't do the trick because silver chloride is insoluble in water. The earliest photographic fixer consisted simply of a concentrated solution of sodium chloride, ordinary table salt:

While certainly convenient, this reaction does not go very far; only some of the silver chloride is dissolved and the rest remains on the paper. More effective than salt is ammonia:

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If you are looking for a household chemical to use as a photographic fixer, clear household ammonia will work better than salt. But far more effective is a specialty chemical which has been synonymous with photographic fixer since 1839: sodium hyposulfite, known to photographers as "hypo," or, by its modern chemical name as sodium thiosulfate:

By rinsing the sensitized paper with sodium thiosulfate, light-sensitive silver chloride goes into solution leaving the unexposed areas of the print white and the exposed areas black.

The image so far is a negative image. To produce a positive print, the process is repeated with a second piece of sensitized paper. The negative is placed on top of this second piece of paper and exposed to sunlight. The black areas of the negative mask the sunlight and the underlying areas of the print remain white. The white areas of the negative allow sunlight to penetrate and the underlying areas of the print turn black. Once fixed, the print bears a positive image, one in which white and black are as they should be.

The silver chloride/albumin combination produces what is known as a "Printing Out Paper," one in which the image becomes visible as the exposure proceeds. Such papers are slow, requiring minutes of exposure to bright light for the image to form. In a "Developing Out Paper," the binder used to form the emulsion is not a reducing agent, giving the unexposed paper a longer shelf life. When such a paper is exposed to light, no image is immediately apparent. The image forms when the exposed paper is placed in a developing bath, a solution of a relatively strong reducing agent. Exposed areas of the print are quickly reduced by the developer, turning black. Before the unexposed areas are reduced, the print is placed in a "stop bath," which destroys any developer remaining on the print. The print is then fixed in the usual fashion.

Photography continued to develop in the twentieth century. Color photography came into its own with the introduction of Kodachrome in 1935. In 1948 Edwin Land introduced a film which incorporated pouches of developer and fixer. When pulled from his "Polaroid" camera, the packets were squeezed onto the film resulting in a self-developing print; color Polaroid film was introduced in 1962. The twenty-first century has seen the maturation of the electronic, film-less camera. In fact, all of the photographs for this book were taken electronically. But the vast majority of photographs at the time of writing continue to be based on the fundamental principles of silver reduction first explored by Herschel, Daguerre, and Talbot.

Material Safety

Locate MSDS's for silver nitrate (CAS 7761-88-8), sodium chloride (CAS 7647-14-5), sodium carbonate (CAS 497-19-8), acetic acid (CAS 64-19-7),

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sodium thiosulfate (CAS 10102-17-7), and household ammonia (CAS 1336-21-6). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[1]

The most likely exposure is to silver nitrate solution, which will blacken the skin. Wash any affected areas immediately with plenty of running water. If eye contact occurs, flush with water and call an ambulance.

You should wear safety glasses and rubber gloves while working on this project. Leftover silver nitrate solution should be collected for recycling or disposal. Leftover egg whites may be thrown in the trash. Spent sodium thiosulfate or ammonia should be washed down the drain with plenty of water unless prohibited by law.

Research and Development

So there you are, studying for a test, and you wonder what will be on it. Study the meanings of all of the words that are important enough to be

included in the index or glossary. You should know all of the Research and Development points from

Chapter 17 and Chapter 19. You should know what substance comprises the black areas of a black

and white photograph; you should know what substance comprises the white areas of a black and white photograph.

You should know three purposes of albumin in the photographic emulsion described in this chapter. Is albumin used in modern photographic emulsions?

You should know the purpose of a photographic fixer; you should know the reactions of sodium thiosulfate and ammonia with silver chloride.

You should know the hazardous properties of sodium chloride, sodium carbonate, silver nitrate, ammonia, and sodium thiosulfate.

You should know the contributions of Daguerre, Talbot, Herschel, and Eastman to the development of photography.

Notes

[1] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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23.3. The creation of albumin photographs consists of five steps; mixing of the albumin/salt solution, coating the paper with albumin/salt solution, sensitizing the paper with silver nitrate solution, exposure through a negative, and fixing with sodium thiosulfate. While just about any paper will work, including handmade paper from Chapter 14, I recommend a good-quality, smooth-textured watercolor paper.[1] You will also need a dozen raw eggs, some table salt (sodium chloride), white household vinegar, silver nitrate, sodium carbonate, and sodium thiosulfate.

To make the albumin/salt solution, you need to separate the yolks from the eggs. One egg at a time, separate the yolk from the whites using your favorite method and place the egg white into a cup. If any egg white is contaminated by yolk, discard the whole egg; egg yolks contain sulfur, which will degrade the quality of your print. Examine each egg white and remove any dirt or bits of eggshell. Once an egg white is free of debris, pour it from the cup into a one-pint glass jar. Continue separating eggs until your jar contains a dozen clean egg whites. To your jar of egg whites add 12 mL of white household vinegar and 12 g of sodium chloride. Put the lid on the jar and shake it until a good head of froth develops. Allow the froth to rise to the surface and then remove it with a spoon, as shown in Figure 23-1(R), leaving clean, clear emulsion in the jar. The emulsion may be prepared in advance and stored in the refrigerator indefinitely. If the emulsion becomes dirty after repeated use, simply shake it and spoon off the froth; most of the impurities will be caught in the froth. A dozen eggs produce enough emulsion for very many prints.

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Figure 23-1. Making the Albumin/Salt Emulsion

When you are ready to coat your paper, remove your albumin solution from the refrigerator and allow it to warm to room temperature. Cut your paper to a convenient size and fold up about 1 cm along each edge of the paper. The effect is that you are forming your paper into a shallow dish. You will need to fold the corners in so that all four edges stand up; do not cut the corners; your paper dish needs to be water-tight. Using a medicine dropper or transfer pipette, withdraw some emulsion from your jar. By putting the tip of the pipette well below the surface, you can extract very clean emulsion, free of bubbles and froth. Deliver the emulsion to your paper dish and continue adding emulsion until the bottom of your dish is covered. You can rock your paper dish from side to side and move emulsion to dry areas with your pipette. The goal is to completely cover the paper with clean, bubble-free emulsion. Once the bottom of your paper dish is completely covered with emulsion, allow it to sit for 3 minutes. At the end this time, use you pipette to return the remaining emulsion to the jar.[2] Then fill your pipette with fresh emulsion and use it to wash any bubbles or dirt from your paper, as shown in Figure 23-2(R). Allow your coated paper to dry overnight; you may lean it vertically against a wall, stand it up in a lab drawer or cabinet, or hang it from a line with clothes pins. Coated paper may be stored indefinitely.[3]

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Figure 23-2. Coating the Paper with Emulsion

Paper should be sensitized in dim light; a 15-watt incandescent bulb or darkroom safe-light may be used. To sensitize your paper you will coat the bottom of your paper dish with silver nitrate solution using a medicine dropper or transfer pipette. However, you should use a different pipette from the one you used for the emulsion to avoid cross-contamination. You should also wear rubber gloves if you don't want your hands blacker than charcoal. The sensitizing solution consists of 30 g of silver nitrate dissolved in 250 mL of water. It may be stored in a brown glass bottle. Paper should not be sensitized, however, until you are ready to use it. To sensitize your paper, pipette enough silver nitrate solution into your coated paper dish to completely cover the bottom. Rock the dish from side to side and use your pipette to move the silver nitrate solution onto any dry areas. Once the bottom of your paper dish is completely covered with silver nitrate solution, allow it to sit for 1 minute. At the end of this time, use your pipette to return the remaining solution to the bottle, as shown in Figure 23-3(R). Stand your sensitized paper vertically against a wall or hang it from a line to dry in the dark for 1 hour. While you are waiting for your sensitized paper to dry, you may coat more paper with emulsion.

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Figure 23-3. Sensitizing the Emulsion

You will need a negative through which to make your exposure. Since you will be making a contact print, the finished print will be the same size as your negative. While you may use a 4x5 inch photographic negative, you may also print a digital image onto transparency film or even onto paper. Most image-manipulation computer programs have an option to produce a negative image. A high-contrast image works best, that is, one with very black blacks and very white whites.

You will also need a printing frame to make your exposure; a standard picture frame works well. The negative and the sensitized paper must be loaded into the printing frame in dim light, the same light you used for sensitizing your paper. Place the negative against the glass of the picture frame. Flatten out your sensitized paper dish and place the sensitized side of the paper against the negative. Then assemble the back of the picture frame so that the negative is held tightly against the sensitized paper so that it holds the sensitized paper tightly against the negative.

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Figure 23-4. Exposing the Print

Figure 23-5. The Finished Print

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You are now ready to make your exposure. Take your printing frame out into bright sunlight. On a cloudy day, bright artificial light may be used. The frame may be set under a bright lamp or it may be placed face-down on an overhead projector. As the exposure proceeds, the areas of your print not masked by the negative will start to darken, as shown in Figure 23-4(R). Particularly watch the exposed area around the edge of the negative. When these exposed areas turn a rich chocolate brown, the exposure is complete. This may take only a few minutes in bright sunlight or as long as an hour under reduced lighting conditions.

Returning to dim light, remove the paper from the printing frame. You should now see an image on the paper; the areas masked by the negative will still be white, while the exposed areas will be dark brown or black. If you were to turn on the lights now, the entire print would quickly turn black. Soak your print for 5 minutes in plain water and then 5 minutes in fixing solution. The fixing solution consists of 2 g of sodium carbonate and 150 g of sodium thiosulfate (aka sodium hyposulfite) in 1 L of water. Whereas only one side was treated with albumin and sensitizer, the whole print needs to be immersed in the fixing bath; you need to remove any silver which may have soaked into the paper. Finally, wash your print for 30 minutes in running water and hang your print on a line to dry. If you reinforce the folds you made at the beginning, the paper will be less likely to curl as it dries.

With this brief introduction to photographic processes, a whole world of alternative processes opens up to you. You may be interested, for example, in the cyanotype, or blueprint, an iron-based photographic process pioneered by none other than John Herschel. You may be interested in the gum-bichromate process, a color process based on

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the interaction of chromium compounds with ordinary water colors. These alternative processes have experienced a boom among artistic photographers in recent years. If you would like to experiment with them, see, for example, References [84] and [82].

Quality Assurance

Include safety information in your notebook on the materials you used in this project. Outline the steps you used and tape your finished photograph into your notebook. A true photograph will sport a sharp, well-defined, and permanent image.

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Notes

[1] I use Canson Montval 90 lb. Watercolor Paper. Many papers will work but the best results are to be had from paper that is heavy, smooth, and well-sized. I typically cut paper to 5x7 inches as a matter of convenience.

[2] It is possible to reuse all three solutions used in this project—the emulsion, the sensitizer, and the fixer—provided that they are not mixed with one another. The emulsion doesn't go bad if refrigerated, it just gets used up. The sensitizer may become depleted after prolonged use; when prints fail to darken in sunlight, it is time to replace the sensitizer. The fixer will also weaken with prolonged use; when fixed photographs darken with time, it is time to replace the fixer.Spent sensitizer and fixer may be recycled to reclaim their silver content. If you have no access to a commercial recycler, these solutions may be evaporated to dryness and then smelted by the procedures of Chapter 9 as if they were ores of silver.

[3] In a classroom situation, there are several ways to work around this overnight drying period. First, you may coat your paper with emulsion one week, sensitize and expose it the following week. Second, you may use paper coated by another student the week before and then coat paper for future students while you are waiting for your sensitized paper to dry. Finally, your instructor may have made up coated papers in advance. Ask your instructor which option is best for you.

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Chapter 24. Solvay (Ammonia)

The greater part of the soda now employed is obtained by Leblanc's process, which while it admits of lixiviating the soda readily and completely, is defective, inasmuch as the residue, or waste as it is technically called, contains nearly all the sulphur used in the manufacture; and that this is not a slight loss may be inferred from Oppenheim's statement, that in the alkali works at Dieuze, Lorraine, the accumulated waste contains an amount of sulfur valued at ?150,000. For every ton of alkali made there is accumulated 1? tons of waste, containing 80 per cent. of the sulphur used in the manufacture; and this waste, until lately thrown on a refuse heap in some fields adjacent to the works, often proved a nuisance in hot weather, giving rise to fumes of suphuretted hydrogen.

— Wagner's Chemical Technology, 1872 AD [1]

24.1. Four score and seven years ago our fathers brought forth on the European continent, a

new industry, conceived in Liberty, and dedicated to the proposition that all men are created equal. The chemical revolution begun by Leblanc in 1790 had passed from France to Britain, where the repeal of the salt tax had stimulated the growth of a booming soda industry. This industry used as raw materials, salt, sulfur, limestone, saltpeter, coal, air, and water; its products were the alkalis, sodium carbonate and sodium hydroxide. Cheap alkalis brought to the ordinary citizen those luxuries which had formerly been enjoyed only by the rich and powerful: glass for bringing light into dark places, paper for bringing the printed word into proletarian homes, and soap for bringing sanitation into cities oppressed by filth and disease. And while the Leblanc process delivered on many of the promises of the revolution, you may very well wonder why the world has little noted, nor long remembered the work which it so nobly advanced. I will tell you.

You will, of course, recall that in addition to valuable alkalis, the Leblanc process produced two waste products, hydrogen chloride and calcium sulfide. Acidic hydrogen chloride gas was sent up the chimney, after which it decimated vegetation in the vicinity of an alkali works. Insoluble calcium sulfide was conveniently disposed of in heaps where the vegetation used to be. Unfortunately, when calcium sulfide reacts with rain water it farts out noxious hydrogen sulfide, a gas which will never have a rose named after it. Not surprisingly, there was little demand for posh plant-side homes in the neighborhood of the Widnes Alkali Works.

As a consequence, alkali manufacturers became popular targets for lawsuits and government regulations. The British Alkali Act of 1863, for example, required the absorption of 95% of the hydrogen chloride produced by the salt cake furnace.[2] This was easily accomplished, hydrogen chloride being quite soluble in water; the waste gas was sent up through a stone tower filled with coke; water dribbling down through the

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tower absorbed the hydrogen chloride, producing aqueous hydrochloric acid. Voila! Clean air. There being little demand for hydrochloric acid, however, it was natural to dispose of it in the nearest river, a loophole closed by the Alkali Act of 1874.

In addition to hydrogen chloride, calcium sulfide was a thorn in the side of soda manufacturers. While the community complained about stinking heaps of tank waste, the manufacturers themselves mourned the loss of valuable sulfur. Sulfur was then imported almost exclusively from Sicily and a monopoly established there in 1838 increased prices sharply.[3] An alternative to sulfur was found in sulfur dioxide, roasted from British and Spanish pyrites, but there was no getting around the stubborn fact that every ounce of sulfur procured from whatever source was destined to wind up as calcium sulfide waste. The one silver lining of the calcium sulfide cloud was that it could be converted into sodium thiosulfate, used by photographers to fix photographs. Even so, the demand for photographic fixer was dwarfed by the supply of tank waste; the noxious heaps became mountains, testing whether any industry so conceived could long endure.

It was into such a world that I, Ernest Solvay, was born in 1838. Son of a Belgian salt-maker, I cut my teeth on salt and at the age of 21 went to work for my uncle, who managed a coal-gas plant. Growing up as I did surrounded by salt, coal, gas, and ammonia, it was natural that I should apply my little gray cells to the problems of the soda industry. Many commercial enterprises had attempted to exploit the reaction of salt, ammonia, and carbon dioxide to produce ammonium chloride and sodium carbonate. But if the ammonia were not recovered, the process would suffer from the waste of ammonia, as the Leblanc process had from the waste of sulfur. To this end, I devised a recovery process using lime to liberate ammonia gas from ammonium chloride. I realized, however, that winning this battle would not be enough to win the war.

Ward and Roebuck, after all, had brought the chamber acid industry into being, but had failed to keep their secrets. Leblanc himself had given birth to the soda industry and yet political upheaval had robbed him of the fruits of his labors. Since it is not in the interest of secrets to be kept, I was careful to patent my ideas both in Belgium and in Britain. I then established a soda works in Belgium which began production in 1864. While the construction of a Solvay plant was more expensive than a comparable Leblanc plant, it required fewer raw materials; thus capital investment was higher but operating costs were lower. I built a second plant in France, but the real key to my commercial success was the sale of licenses abroad; my particular un-kept secrets were going to pay as they went. Solvay plants were established by Brunner, Mond & Co. in Britain and by Solvay Process Co. in the United States, paying royalties on every ton of soda produced.

The established alkali manufacturers were not about to surrender without a fight, however. In 1868 Henry Deacon introduced a process for turning waste hydrogen chloride into bleaching powder, which found a ready market in paper and textiles.[4] In 1887 Alexander Chance finally succeeded where so many others had failed in recovering sulfur from tank waste. But these piece-meal improvements would make a patchwork of existing alkali plants, sacrificing simplicity for efficiency. With lower operating costs, I was able to drop the price of soda to the point that my competitors were forced to sell it at

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a loss. Leblanc soda works, in an ironic twist, would be kept afloat only by sales of bleaching powder produced from acid formerly run to waste.

World production of soda in 1863 had been 150,000 tons, all produced in Leblanc plants. By 1902 world production of soda would soar to 1,760,000 tons, over 90% of which would be produced in Solvay plants.[5] It would be altogether fitting and proper to celebrate the victory of cleaner and more efficient processes over those crippled by pollution and waste. But, in a larger sense, we can not dismiss—we can not forget—we can not overlook those who struggled to bring light, literature, and lather to all peoples. It is rather for us to be here dedicated to the great task remaining before us—that from these honored dead we take increased devotion to that cause for which they gave the last full measure of devotion.

Notes

[1] Reference [29], pp. 184-185.[2] Reference [68], p. 89.[3] Reference [75], p.33.[4] Reference [60], pp. 95-98.[5] Reference [52], p. 59.

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24.2. In order to understand the Solvay process, it is necessary to understand the chemistries of carbon dioxide and ammonia. We first encountered carbon dioxide (CO2) way back in Chapter 1 as a product of the combustion of charcoal. We learned in Chapter 4 that yeasts fart carbon dioxide when they consume honey. Chapter 18 characterized carbon dioxide as a mildly acidic gas which reacts with bases to form carbonate or bicarbonate[1] salts. Carbon dioxide may be liberated from these salts either by heating the bejeezus out of them or by reacting them with acids. The properties of carbon dioxide are summarized in Equation 24-1.

Equation 24-1. Reactions Involving Carbon Dioxide

Equation 24-2. Reactions Involving Ammonia

We met ammonia (NH3) in Chapter 12 as a mild alkali suitable for dissolving indigo dye. Ammonia reappeared in Chapter 22 as a by-product of the distillation of coal tar. The "ammonia" sold at the grocery store is actually a solution of ammonia in water. Just as carbon dioxide combines with water to form carbonic acid, ammonia combines with water to form the alkali, ammonium hydroxide (NH4OH). Ammonium hydroxide participates in the usual acid-base reactions to form ammonium salts. Ammonia may be liberated from these salts either by heating the bejeezus out of them or by reacting them with stronger alkalis like sodium or calcium hydroxide. The properties of ammonia are compared to those of carbon dioxide in Equation 24-2.

Carbonic acid reacts with a base to form a bicarbonate salt and water; ammonium hydroxide reacts with an acid to form an ammonium salt and water. It is only natural that carbonic acid should react with ammonium hydroxide to produce ammonium bicarbonate and water:

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H2CO3(aq) + NH4OH(aq) = NH4HCO3(aq) + H2O(l)

When a solution of ammonium bicarbonate evaporates, it leaves behind solid ammonium bicarbonate, a salt with the properties of both bicarbonate salts and ammonium salts. Like other bicarbonate salts, heating the bejeezus out of it liberates carbon dioxide; like other ammonium salts, heating the bejeezus out of it liberates ammonia. The curious thing about ammonium bicarbonate it that when you liberate bejeesical carbon dioxide and bejeesical ammonia, the only thing left is bejeesical water. Thus ammonium bicarbonate is entirely bejeesical in nature; heat it and it just disappears. This property makes it a useful ingredient of baking powders because in the oven it is able to leaven breads and cakes without leaving anything behind.

Ammonium bicarbonate is crucial to the Solvay process because it undergoes a metathesis reaction with sodium chloride to produce sodium bicarbonate and ammonium chloride:

NaCl(aq) + NH4HCO3(aq) = NaHCO3(s) + NH4Cl(aq)

You are probably thinking that sodium bicarbonate is soluble in water. That is so, but it is less soluble than sodium chloride, ammonium bicarbonate, or ammonium chloride. Thus in a very concentrated solution, sodium bicarbonate will be the first to precipitate. This metathesis reaction is the key to the Solvay process.

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Figure 24-1. The Solvay Soda Process

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The three critical reactions of the Solvay process occur simultaneously in reactor (e) of Figure 24-1. Remember that in a process schematic, reactants enter a reactor from the left and products leave to the right. An aqueous solution of ammonium hydroxide and sodium chloride enters reactor (e), an absorber, where it meets a stream of carbon dioxide gas. The carbon dioxide dissolves to produce carbonic acid, which reacts with ammonium hydroxide to produce ammonium bicarbonate. If there were plenty of water, everything would remain in solution—sodium ion, chloride ion, ammonium ion, and bicarbonate ion. But if water is limited, that is, if the solutions are concentrated, then sodium bicarbonate precipitates and ammonium chloride remains in solution. The liquid returns to reactor (c) and the solid goes to furnace (f), where bejeesical carbon dioxide is driven off. The sodium carbonate which remains is the intended product of the whole process.

Reactors (e) and (f) are the stars of the Solvay process; all of the other reactors play supporting roles. Absorber (e) requires ammonium hydroxide from another absorber, (d). Absorber (d) requires gaseous ammonia from a still, (c). Still (c) requires ammonium chloride recycled from absorber (e) and slaked lime from a slaker, (b). Finally, slaker (b) requires lime from a furnace, (a). As raw materials, the whole process requires limestone, water, and salt. It produces waste calcium chloride and a single product, sodium carbonate. If you add up all of the reactions from all of the reactors, canceling compounds that appear on both sides of the equal sign, the result is remarkably simple:

CaCO3(s) + 2 NaCl(aq) = CaCl2(aq) + Na2CO3(s).

This reaction may seem familiar to you, since it was the second reaction discussed in Section 7.2. Our solubility rules tell us, however, that this reaction goes the other way. Soluble calcium chloride and sodium carbonate will react in aqueous solution to produce insoluble calcium carbonate and soluble sodium chloride. The Solvay process, with its elegant scheme for recycling carbon dioxide and ammonia, is simply an elaborate means for pushing this classic metathesis reaction backwards.

Compare the Solvay process to the Leblanc process, Figure 20-2, and you will immediately see why the one superseded the other. Both processes require limestone, salt, and water as raw materials, but while the Leblanc process additionally requires sulfuric acid, the Solvay process recycles its ammonia. Both processes produce the same product, sodium carbonate, but while the Leblanc process produces two hazardous waste products, the Solvay process produces only one relatively harmless waste product. In fact, calcium chloride is useful as a road de-icer, though the market is not large enough to absorb the huge amounts produced each year. So while the waste produced by a Solvay plant is not nearly as hazardous as those produced by a Leblanc plant, calcium chloride remains a disposal problem for the soda industry.

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Material Safety

Locate MSDS's for ammonia (CAS 7664-41-7), carbon dioxide (CAS 124-38-9), and ammonium bicarbonate (CAS 1066-33-7). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[2]

Your most likely exposure will be to ammonia fumes. If a persistent cough develops, seek medical attention.

You should wear safety glasses while working on this project. All activities should be performed in a fume hood or with adequate ventilation. Leftover materials can be flushed down the drain with plenty of water.

Research and Development

You are probably wondering what will be on the quiz. You should know the meanings of all of the words important enough to

be included in the index or glossary. You should know the Research and Development points from Chapter

18 and Chapter 20. You should know the reactions of Figure 24-1. You should know the hazardous properties of ammonia, carbon

dioxide, and ammonium bicarbonate. You should be able to reproduce Figure 24-1 and to explain this

schematic in your own words. You should be able to explain the advantages of the Solvay process

over the Leblanc process.

You should be able to tell the story of Earnest Solvay.

Notes

[1] The modern name for the bicarbonate ion is the hydrogen carbonate ion. The older name, however, continues to be widely used.

[2] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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24.3. It rises from the Earth to Heaven and descends again to the Earth, and receives power from Above and from Below. Thus thou wilt have the glory of the Whole World. All obscurity shall be clear to thee.

— The Emerald Tablet of Hermes Trismegistos

A lime kiln, a slaker, a still, two absorbers, and a furnace make up the complete Solvay process. None of this is new; you have known how to make lime since Chapter 10; you grasped the secrets of distillation in Chapter 16; you learned how to build an absorber in Chapter 18; and the final furnace is quite similar to the one described in Chapter 20. You could build the whole plant out of pottery and soft-drink bottles if you wanted too, but I am afraid that few readers will want to go to all that trouble to produce something that might have been extracted in one step from seaweed ashes. After batteries, dyes, and photographs, the prospect of making one white powder from another might be a little bit anti-climactic. So how about a seance? Let us conjure some spirits, materialize some ghosts, bottle some genies, herd some bejeesical cats. Let us make some ammonium bicarbonate. Okay, so it is not going to be the perfect birthday present for Mom. You are not going to save a bottle of it for graduation day or mount it on your key-chain. Look, nobody is holding a gun to your head; if the prospect of making something "rise from Earth to Heaven and descend again to the Earth" does not float your boat, then go directly to the next chapter; do not pass "Go," do not collect the glory of the Whole World.

For anyone who made it past the last sentence, let us see whether we can clarify some obscurity. You have been using ammonia since Chapter 12, whether home-grown or store-bought. You are undoubtedly familiar with ammonia's unmistakable aroma, but how does it get from the bottle to your nose? The stuff in the bottle is an aqueous solution of ammonium hydroxide, NH4OH. Escaping from the bottle is bejeesical ammonia, NH3, the gas which makes your eyes water. Unfortunately, the word ammonia is often used indiscriminately to describe both the gas and its solution. Commercial ammonium hydroxide is 28-30% ammonia in water; household "ammonia" is more dilute. For this project we need gaseous ammonia, which we can get by distilling household ammonia.

In addition to ammonia, we need some carbon dioxide. In the industrial process this comes from the lime kiln, but we can use the carbon dioxide collected in Section 4.3 for this very purpose. Figure 4-3 details the fermentation lock consisting of a balloon glued to a bottle cap with a hole in it. This balloon is either still screwed onto your mead or you have wrapped its neck around and pencil and secured it with a twist tie. You will need a minimum of 2 L of carbon dioxide. If your balloon is round, measure its circumference with a piece of string; the circumference should be at least 50 cm. If your balloon is oblong it should be at least as big as a 2-liter soft-drink bottle. With your balloon of carbon dioxide handy, it is time to distill your ammonia.

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Figure 24-2. Distillation and Absorption

Set up your still just as you did in Chapter 16, but instead of filling your Erlenmeyer flask with mead, fill it with 400 mL of household ammonia and a few boiling chips. For the receiver you can use our old friend, the 2-liter soft-drink bottle. Since we are collecting a gas there is no need to cool the receiver; there is no need for ice. Slip the receiver over the condenser tube and place the still on a hot-plate. Set the hot-plate power to "high," but as soon as the pot begins to boil turn the power down to half. Adjust the hot-plate power to keep the pot boiling without taking the head temperature above 82?C (180?F); you want to distill gaseous ammonia, not water. If you are not paying attention, the head will climb too high, too fast, and you will have to start over. If, however, you separate the gas by fire, the fine from the gross, gently and with great skill, then the head temperature will climb gradually to 82?C. You will notice what looks like steam condensing on the sides of the receiver; this is concentrated ammonium hydroxide. Allow a few mL to collect in the bottom of the receiver and then shut down the distillation. Remove the receiver from the still, drain the ammonium hydroxide into a small beaker, and using a graduated pipette or graduated cylinder pour 2 mL of ammonium hydroxide back into the receiver. The reason for returning this ammonium hydroxide will be explained shortly.

Once the ammonia distillation is complete, transfer the balloon from your mead to your receiver as quickly as possible. Pinch the neck of your balloon, unscrew it from your mead, screw it onto the receiver, and release the neck of the balloon. Over the course of a few minutes bejeesical carbon dioxide will be absorbed by the ammonia; your receiver has miraculously changed into an absorber. The balloon will shrink as carbon dioxide

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becomes carbonic acid, as shown in Figure 24-2(R). The bottle will get hot as carbonic acid reacts with ammonium hydroxide to form ammonium bicarbonate. Roll the bottle from side to side to keep the walls wet. You will notice what looks like condensation on the inside of the bottle, but as the reaction proceeds this condensation will sprout crystals, as shown in Figure 24-3(L). While your crystals are growing, let us take a look at the stoichiometry.

CO2(g) + NH3(g) + H2O(l) = NH4HCO3(s)

At room temperature and normal atmospheric pressure a mole of gas occupies 24.5 liters. This unit factor, (24.5 L/mol CO2) or (24.5 L/mol NH3) allows us to answer stoichiometric questions:

Q: How many grams of ammonium bicarbonate can be produced from the reaction of 3.0 L of carbon dioxide with excess ammonia and water?Q: How many grams of ammonium bicarbonate can be produced from the reaction of 2.0 L of ammonia with excess carbon dioxide and water?Q: How many grams of ammonium bicarbonate can be produced from the reaction of 2.0 g of water with excess ammonia and carbon dioxide?

Q: How many grams of ammonium bicarbonate can be produced from the reaction of 3.0 L of carbon dioxide with excess ammonia and water?

A: 9.6 g.

Q: How many grams of ammonium bicarbonate can be produced from the reaction of 2.0 L of ammonia with excess carbon dioxide and water?

A: 6.4 g.

Q: How many grams of ammonium bicarbonate can be produced from the reaction of 2.0 g of water with excess ammonia and carbon dioxide?

A: 8.8 g.

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Figure 24-3. Spirit Made Flesh

So what is the theoretical yield, 22.6 grams? No, to produce 6.4 g of ammonium bicarbonate you need 2.0 L of carbon dioxide and 2.0 L of ammonia and 1.5 g of water. If you have 2.0 L of ammonia you can make at most 6.4 g of ammonium bicarbonate, no matter how much carbon dioxide and water are present. In this case the limiting reagent, the one which runs out first, is ammonia and the theoretical yield is 6.4 grams. If your balloon holds less than 2.0 L of carbon dioxide it will completely collapse; in this case carbon dioxide is the limiting reagent and your theoretical yield will be somewhat less than 6.4 g. Similarly, if there is less than 1.5 g of water present, water is the limiting reagent. We added 2 mL of ammonium hydroxide (ammonia in water) to ensure that there would be enough water present. Adding more ammonium hydroxide might have increased our yield since it supplies both water and ammonia, but we would have wound up with a solution of sodium bicarbonate rather than crystals. Limiting the water allowed us to collect crystals immediately without having to wait for the excess water to evaporate.

When the balloon stops shrinking and the bottle is cool, remove the balloon and cut the bottle open at the shoulder. Use a spoon or spatula to push the ammonium bicarbonate crystals to the bottom of the bottle. Trim away the walls of the bottle as you go until you are left with a shallow dish filled with crystals, as shown in Figure 24-3(R). Allow your crystals to dry overnight and then weigh them to get your actual yield, which is often expressed as a percentage of the theoretical yield.

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Consider what has happened here; you started with an empty bottle; everything that went into the bottle was a gas, yet from these bejeesical spirits a solid body emerged. As an infant it smells of ammonia but when you allow it to dry overnight the smell almost goes away. It has so much potential; it might make bread or cake or soda, but this particular body is destined for a trial by fire. Place it into a small beaker and heat it with a hot-plate or spirit lamp and as it gets old and decrepit it will begin to smell of ammonia once more. It begins to lose weight and finally just fades away, departed once more into the spirit world. In the end, the only thing left of it is a story.

Quality Assurance

Include in your notebook safety information for ammonia, carbon dioxide, and ammonium bicarbonate. Record the actual yield of ammonium bicarbonate expressed as a percentage of the theoretical yield. Include a photograph of your crystals to illustrate your telling of their poignant saga.

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Chapter 25. Dow (Electrochemicals)

[1]

25.1. (1890) For his undergraduate chemistry project Herbert Dow had identified Michigan

brines which were unusually high in bromine. Bromine was used to produce patent medicines and photographic chemicals; the only commercial source was Die Deutsche Bromkonvention, a German cartel. Dow formed the Midland Chemical Company in 1890 to produce electrolytic bromine from brine, but when he wanted to expand into chlorineand caustic soda (chlor-alkali), his financial backers balked. Dow left in 1895 to form the Dow Chemical Company, which absorbed Midland in 1914. Though it survived competition from German bromine and British bleach, Dow remained a small company until the First World War interrupted trade with Germany. The war made it possible to move into aniline dyes and pharmaceuticals, areas dominated by the German giants, and Dow began manufacturing phenol, aspirin, and synthetic indigo. Its expertise in bromine chemistry led it to develop anti-knock gasoline additives in partnership with Ethyl Corporation. Its roots in electrolysis allowed it to extract magnesium metal from sea-water by the end of the Second World War. Its increasing proficiency with organic chemicals led it to develop silicone sealants in partnership with Corning Glass. Its early work with chlorinated fungicides led it to develop defoliants for the American military during the Vietnam conflict. Its early success with transparent polystyrene led it to develop Saran Wrap and Ziploc bags. Following its acquisition of Union Carbide, Dow Chemical would begin the twenty-first century as the largest producer of chemicals in the world.[2]

(1888) Two years before Dow formed Midland, American-born Hamilton Castner went to work for the British Aluminum Company. Aluminum manufacture at that time required elemental sodium and Castner had invented a process for making sodium from soda. Unfortunately for him, the Hall-H 鲯 ult process for making aluminum electrolytically was about to eliminate demand for sodium in aluminum manufacture. Desperately seeking a more efficient route to sodium, he began electrolyzing brine with a mercury cathode. Though the process never produced sodium cheaply enough to compete in aluminum manufacture, it produced chlorine and very pure caustic soda. It was, in fact,

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for the sake of caustic soda that the Austrian paper pulper, Karl Kellner, patented an almost identical process in 1894. The two Athanors resolved their initial disputes and formed the Castner-Kellner Alkali Company, which took over the British Aluminum Company in 1897.[3] Castner-Kellner would, in turn, be swallowed by Brunner Mond, which manufactured soda under license from Solvay. The availability of cheap electrolytic chlorine and caustic soda was the final nail in the coffin of the Leblanc soda process. Brunner Mond would form the core of the industrial giant, Imperial Chemical Industries (ICI), which would begin the twenty-first century as the tenth largest chemical company in the world.

(1888) When Castner went to work for British Aluminum, aluminum was a precious metal. The third most abundant element in the Earth's crust, aluminum could not be won from its ore by smelting, carbon being an insufficiently strong reducing agent. Elemental sodium was the industrial reducing agent of choice for winning aluminum from its oxide, alumina, and the high price of sodium translated into high prices for aluminum. While sodium could be produced from the electrolysis of molten sodium hydroxide, the melting point of alumina, 2051?C, ruled out the electrolysis of molten alumina as a viable industrial process. An American student, Charles Martin Hall, learned of these difficulties in his undergraduate chemistry classes at Oberlin and set out to find a flux which would lower the melting point of alumina to an attainable level. The mineral cryolite proved to be such a flux, lowering the melting point to a mere 950?C, well within the range of even modest furnaces. Using home-made batteries he produced his first electrolytic aluminum in 1886. Hall attempted to commercialize an electrolytic process in collaboration with the Cowles Electric Smelting and Aluminum Company, developers of the electric smelting furnace.[4] Meanwhile in France another college student, Paul H鲯ult, had independently found cryolite to be an acceptable flux for alumina and had begun using a small dynamo to produce electrolytic aluminum. It is in the interest of mortals to keep secrets and consequently Hall, H 鲯 ult, and Cowles engaged each other in lengthy

legal battles. But it is not in the interest of secrets to be kept; H 鲯 ult would license his process to the British Aluminum Company and the Hall process would be adopted by the Pittsburgh Reduction Company, destined to become the Aluminum Company of America (AlCoA).

(1881) Seven years before Hall obtained his first aluminum, Thomas Alva Edison opened the first electrical power station at Pearl Street in New York City. Davy had invented the electric arc light in 1808 but the power available from batteries was insufficient for commercial application. Oersted had demonstrated electromagnetism in 1820 and Faraday had discovered electromagnetic induction in 1831, but it would fall to Werner von Siemens in 1867 to develop the dynamo, a device for converting mechanical power into electrical power. With increasing demand for electric lighting came the widespread availability of electrical power for other uses, including electrolytic production of chlorine, caustic soda, and aluminum. The early electrochemical industry grew up near sources of cheap hydroelectric power, for example, at Niagara Falls. The industry became less centralized as the AC system of Nicola Tesla and George Westinghouse made it possible to distribute electricity from remote power plants.

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Edison's General Electric would dwarf even the largest chemical company at the beginning of the twenty-first century, with total revenues almost five times that of Dow; in chemical sales alone it would rank as sixteenth largest chemical company in the world.

(1864) Seventeen years before the electric light, the introduction of the Solvay process had signaled the beginning of the end for the Leblanc soda process, then the central process in chemical industry. In the face of increasing opposition to its practice of dumping waste hydrogen chloride into the environment, the industry found that hydrogen chloride reacted with manganese dioxide to produce chlorine. Chlorine was further reacted with lime to produce calcium hypochlorite, known in the trade as "bleaching powder." As the Solvay process was producing soda at ever lower prices, the Leblanc industry came to rely on sales of bleaching powder to the textile and paper industries as its major source of revenue. This last refuge of the huge Leblanc industry would be taken from it by Castner and Kellner's process for electrolytic chlorine; the Leblanc process would become extinct after the First World War.

(1834) Thirty years before the first Solvay soda plant Davy's protege, Michael Faraday, first began to understand the nature of electrolysis. An electrolysis cell was simply a voltaic cell with the poles reversed; in a voltaic cell the anode, the site of oxidation, was the negative pole and the cathode, the site of reduction, was the positive pole; in the electrolysis cell the anode was forced to be positive and the cathode negative. Just as a mole of a substance can be delivered by weighing out one formula weight of that substance, a mole of electricity can be delivered by passing a current of 96,500 amps through a solution for 1 second; 96.5 amps for 100 seconds; or 9.65 amps for 1000 seconds. The quantity 96,500 amp-sec was, in essence, the formula "weight" for electricity. This new unit, the faraday, allowed stoichiometric calculations to relate the weights of reactants and products to the amount of electricity produced by a battery or consumed by an electrolytic cell.

(1825) Nine years before Faraday demonstrated the chemical equivalence of electricity, Hans Christian Oersted reacted an amalgam of potassium and mercury with aluminum chloride. When the mercury was distilled, it left behind a new element, aluminum. This remarkable metal had a low density and resisted corrosion to a degree comparable to that of gold and silver. Aluminum began to be produced commercially and might have been extremely useful had it not been so expensive. Napoleon III, for example, reserved his aluminum tableware for honored guests on special occasions. Aluminum would remain a precious metal for most of the nineteenth century, its price falling below that of silver only after the introduction of the Hall-H鲯 ult process.

(1797) After millennia of wandering I passed from Lavoisier's Elements of Chemistry to the seventeen year old son of a Penzance wood carver. Forced to support his family by the death of his father three years earlier, this Humphry Davy had apprenticed himself to a saw-bones, but the spark burned too brightly within him to settle for a life of mere quackery. In 1798 he went to work for Thomas Beddoes investigating the therapeutic use of nitrous oxide (laughing gas) and his reputation as a chemist became such that he was

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appointed Professor of Chemistry at the Royal Institution in 1802. In 1808 Davy used the recently introduced voltaic battery to decompose molten potash and soda, producing two hitherto unknown elemental metals and naming them potassium and sodium, respectively. In 1810 he argued that the greenish-yellow gas liberated from common salt by manganese dioxide was, in fact, an element and gave it the name, chlorine. He further demonstrated that chlorine could be produced simply by passing an electrical current through salt water. This process, electrolysis, was to become a powerful tool, turning erstwhile products into reactants, reactants into products, in essence providing a means for driving redox reactions backwards.

Notes

[1] Reference [29], p. 114.[2] Reference [54], July 29, 2002, p. 16.[3] See Reference [61], p. 80.[4] For details of the Cowles-Hall dispute see Reference [56].

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25.2. I was a young teenager when I re-meme-bered the electrolysis of salt water. I was not yet a chemist, but I was keenly interested in science. The chemistry sets of the day did little to encourage an interest in chemistry; they went so far out of their way to make sure that nothing in the set could possibly be dangerous that little remained to interest the inquisitive mind. My scientific passion was inflamed, not by the watered-down kiddie-science of the sixties and seventies, but by the boy-scientist[1] books of the twenties and thirties. Among the admittedly dangerous activities described there, the electrolysis of salt water sparked my interest. With a few common materials I could make a poison gas on the one hand and an explosive gas on the other. Is this a great world, or what?

In Chapter 21 we learned that a voltaic cell, or battery, separates the two half-reactions of a redox reaction so that they take place at two different electrodes; the oxidation takes place at the anode and the reduction at the cathode. On a process schematic the electrons are shown explicitly as if they were reactants and products and these electron "pipes" represent the wires connected to the battery. The wire on the reactant side of the schematic represents the "+" terminal of the battery and the one on the product side represents the "-" terminal. The "desire" of the electrons to cross from anode to cathode is expressed as the electromotive force (EMF), with units of volts. The rate of flow of electrons from anode to cathode is expressed as the current, with units of amps. As a convention, electrons flow from top to bottom in the process schematic of a voltaic cell, as if they were flowing downhill. Reversing the normal direction of electron flow results in an electrolytic cell, one in which electrons are "pumped" from the bottom to the top of the process schematic. Of course, "down" and "up" are merely conventional directions for these process schematics and have nothing to do with the physical geometry of the cell. Let us now explore three applications of electrolysis, the lead-acid storage battery, the Hall-H鲯 ult aluminum process, and the Castner-Kellner chloralkali process.

Figure 25-1 shows a process schematic of the lead-acid storage battery; the anode is lead, the cathode is lead oxide, and the electrolyte is sulfuric acid. If the lead and lead oxide were to come into contact, they would react with one another, electrons flowing from one to the other. But by separating the lead from the lead oxide, we can require the electrons to pass through an external circuit to get from the anode to the cathode. Lead and lead oxide are consumed, lead sulfate is produced, and the sulfuric acid in the electrolyte is converted to water. The astute reader will be able to balance the lead-acid redox reaction after examining the schematic; reactants for the skeleton reaction are on the left of the figure and products are on the right:

? Pb(s) + ? PbO2(s) + ? H2SO4(aq) = ? PbSO4(s) + H2O(l)

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Figure 25-1. The Lead-Acid Cell

In operation, the negative terminal of the battery connects to the lead anode and the positive terminal connects to the lead oxide cathode. But suppose for a moment that we have access to a stronger battery than the lead-acid cell, that is, one with a higher EMF. Further suppose that we connect the two batteries anode-to-anode and cathode-to-cathode. Which direction will electrons flow? Normally, electrons flow out of the anode, but with two anodes connected to one another, the anode of the stronger battery "wins," pushing electrons into the weaker battery and converting its anode into a cathode. Similarly the stronger cathode pulls electrons from the weaker battery, converting its cathode into an anode. This drives the redox reaction of the weaker battery backwards, turning reactants into products and vice versa. In other words, electrolysis recharges the lead-acid cell. In practice, the EMF needed to recharge a car battery comes from the car's alternator rather than from another battery. But no matter what the source of the electrical power, the process schematic of the electrolytic cell is the mirror image of the voltaic one.

A second example of an electrolytic process, the Hall-H 鲯 ult process, has an operating temperature of 950?C. The electrolyte consists of molten alumina, Al2O3, and because the melting point of alumina is 2051?C, a flux is used to lower its melting point. The cryolite flux, Na3AlF6, is not consumed and so does not appear on the process schematic, Figure 25-2. The molten aluminum produced by the process serves as the cathode. Since aluminum is more dense than alumina, it sinks to the bottom of the reactor as a separate liquid phase. The carbon anode is consumed in the process, yielding carbon dioxide gas. As with the lead-acid battery, the reactants and products of the skeleton reaction appear on the schematic and from them the redox reaction may be balanced:

? Al2O3(l) + ? C(s) = ? Al(l) + ? CO2(g)

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Figure 25-2. The Hall-H鲯 ult Process

A third example of electrolysis is provided by the chloralkali cell, shown in Figure 25-3. Similarly to the lead-acid battery, the electrolyte is aqueous, in this case a solution of sodium chloride. The anodic half-reaction is:

2 NaCl(aq) = Cl2(g) + 2 Na+(aq) + 2 e-

Figure 25-3. The Chloralkali Process

The cathodic half-reaction converts water into hydrogen gas and hydroxide ion:

2 Na+(aq) + 2 H2O(l) + 2 e- = H2(g) + 2 NaOH(aq)

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A simple chloralkali cell consists simply of two inert electrodes in a salt solution. But a complication arises because the chlorine gas produced at the anode reacts with sodium hydroxide to produce sodium hypochlorite, NaOCl, the active ingredient in ordinary laundry bleach:

Cl2(g) + 2 NaOH(aq) = NaOCl(aq) + NaCl(aq) + H2O(l)

Though sodium hypochlorite is a useful product, a simple chloralkali cell produces a mixture of aqueous sodium chloride, sodium hydroxide, and sodium hypochlorite. The Castner-Kellner cell was designed to produce pure chlorine and sodium hydroxide by physically separating them from one another. The anodic half-reaction is the same as before, but the Castner-Kellner electrolysis cell employs a metallic mercury[2] cathode, at which the half-reaction is:

2 Na+(aq) + 2 e- = 2 Na(Hg)

The product of this half-reaction is metallic sodium dissolved in metallic mercury. This mercuric solution flows from the electrolysis cell to a separate tank where its dissolved sodium reacts with water:

2 Na(Hg) + 2 H2O(l) = H2(g) + 2 NaOH(aq)

The sum of these last two reactions is the same as our previous cathodic reaction; the Castner-Kellner cell simply ensures that the chlorine and sodium hydroxide are produced in separate containers. Relieved of sodium, valuable mercury is returned to the electrolysis cell and every effort is made to conserve it. Because it is not consumed in the process, mercury does not appear in Figure 25-3.

We must touch on one more topic before moving on to the construction of a simple chloralkali cell. In Chapter 15 we learned to use the formula weight to answer stoichiometric questions. In practice, electrons are measured out, not by weight, but by a combination of current and time. The unit factor, (96,500 amp?sec/mol e -) may be used to answer stoichiometric questions involving voltaic and electrolytic cells. For example:

Q: How long will it take for a chloralkali cell passing a current of 100 mA to produce 500 mL of hydrogen gas?

Q: How long will it take for a chloralkali cell passing a current of 100 mA to produce 500 mL of hydrogen gas?

A: 5.6 hours. Hint: because a higher current would require less time, "100 mA" goes in the bottom. Try starting with (500 mL H2/100 mA) as the first unit factor in the chain.

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Material Safety

Locate MSDS's for chlorine (CAS 7782-50-5), sodium hydroxide (CAS 1310-73-2), sodium chloride (CAS 7647-14-5), sodium hypochlorite (CAS 7681-52-9), and hydrogen (CAS 1333-74-0). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and either the NFPA or HMIS ratings for each material.[3]

Your most likely exposure is to chlorine gas. If breathing becomes difficult get plenty of fresh air and call for an ambulance. Be aware that hydrogen gas is flammable and that sodium hypochlorite will bleach clothing.

You should wear safety glasses while working on this project. All activities should be performed in a fume hood or with adequate ventilation. Leftover solution may be flushed down the drain. The hydrogen produced should be carefully burned. The chlorine should be released outdoors or into a fume hood.

Research and Development

You should not remain ignorant if you are to proceed in the Work. Know the meanings of those words from this chapter worthy of

inclusion in the index or glossary. You should have mastered the Research and Development items of

Chapter 21 and Chapter 22. Know one new unit factor from this chapter and be able to use it to

answer stoichiometric questions involving electrons. Know the hazardous properties of chlorine, sodium hypochlorite,

sodium hydroxide, and hydrogen. Know the balanced redox reactions for the lead-acid cell, the Hall-H鲯

ult cell, and the Castner-Kellner cell. Be able to distinguish between the hazardous properties of mercury,

mercuric chloride, and methyl mercury.

Know the contributions of Dow, Castner, Kellner, Hall, H 鲯 ult, Faraday, and Davy to the understanding of the spark.

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Notes

[1] Such books may also be of interest to girl-scientists who are capable of overlooking the un-apologetic sexism of the era in which they were written.

[2] Mercury has acquired a bad reputation, some of it well-deserved. While metallic mercury is essentially non-toxic and quite insoluble in water, some small fraction of the mercury used in the Castner-Kellner cell will make its way into waste water. Discharged into lakes and streams, bacteria metabolize it into very toxic compounds such as methyl mercury. These compounds make their way up the food chain, being concentrated in the carnivores at the top, such as the tuna. Like the Leblanc soda industry before it, the chloralkali industry has been pressured to address its pollution problems and has gradually replaced aging Castner-Kellner cells with other designs which do not use mercury.

[3] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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25.3. Figure 25-4. Reclaiming Carbon Rods

In the simplest possible terms, we shall pass an electric current through salt water. All that is absolutely necessary is to connect two wires to a battery and dip them into a glass of salt water. Copper wires would rapidly corrode, however, and so we shall use carbon electrodes, just as we did in the aluminum-alkali battery (Chapter 21). Convenient carbon electrodes may be scavenged from non-alkaline flashlight batteries, even "dead" ones. Such "heavy duty" batteries have an outer zinc case and a carbon rod down the center; alkaline have the reverse geometry, with a zinc rod down the center. Wearing safety glasses and gloves, use a pair of pliers to dismantle two standard (Leclanche) flashlight batteries. Peel off the steel jacket, exposing the zinc can underneath. The button end of the battery is usually sealed with a plastic cap; pry this cap off as shown in Figure 25-4(L) to reveal the carbon rod beneath. Carefully twist and pull the carbon rod from the battery. The black stuff in the zinc can is a paste of manganese dioxide and ammonium chloride insulated from the zinc by a paper sleeve. Save the carbon rods and dispose of the rest in the trash.

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Figure 25-5. Constructing the Chloralkali Cell

The body of the cell will be constructed from a 2-liter plastic soft-drink bottle. Use a knife to cut the top from the bottle at the shoulder. Poke two small holes on opposite sides of the bottle near the bottom and slide a carbon rod into each hole, leaving half an inch of each rod outside the bottle. Apply glue[1] to make a water-tight seal between the rods and the bottle, as shown in Figure 25-5(L), and allow the glue to dry or set. You will also need two smaller plastic bottles of such a size that they fit inside the 2-liter bottle, side-by-side. Use a knife to cut the bottoms from these bottles and trim them so that they fit completely inside the 2-liter bottle, one over each of the carbon rods.

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Figure 25-6. Filling the Chloralkali Cell

Fill the 2-liter bottle to the brim with water and add salt (sodium chloride), stirring until no more will dissolve. With their caps removed, push the two smaller bottles into the larger one, bottoms first, so that each one sits over a carbon rod. The mouths of these bottles should be flush with the surface of the saturated salt solution. Screw on the caps, trapping as little air as possible in the small bottles, and lift them gently, as shown in Figure 25-6(R); the water level in the 2-liter bottle will fall but the smaller bottles will remain filled with water.

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Figure 25-7. Operating the Chloralkali Cell

Use a wire to connect each of the carbon rods to a terminal of a lantern battery. Very soon bubbles will begin to form on the surface of the carbon electrodes and will rise to be collected in the two smaller bottles, as shown in Figure 25-7. One of these gases is hydrogen, the other chlorine. If you have been paying attention, you should be able to predict which is which by carefully considering Figure 25-3 and noting which electrode is connected to the positive and which to the negative terminal of the battery. Just picture electrons coming out of the negative terminal and flowing into the positive one. Hydrogen will collect more quickly than chlorine because some of the chlorine will dissolve in the increasingly alkaline solution to form sodium hypochlorite, the active ingredient in ordinary laundry bleach. As the gases displace the water in the smaller bottles, the water level in the larger one will rise; lifting the smaller bottles from time to time will prevent the larger one from overflowing. Nevertheless, it is probably a good idea to keep your chloralkali generator in a sink, tub, or bucket while it is in operation.

When one bottle is full of hydrogen gas, it is time to test your products. Carefully lift the smaller bottles from the larger one and place their open bottoms on a counter top or table which will not be harmed by bleach. The bottle which was full of gas will contain hydrogen. Light a splint or fireplace match and hold it against the bottom of the hydrogen bottle as you turn it over. The hydrogen will light with a satisfying pop. Be careful to point the bottle away from arms, faces, and other flammable materials.

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The other bottle contains chlorine, the toxic gas, used on the battlefields of World War I. It is also the gas used to disinfect drinking water and swimming-pool water. The smell of laundry bleach is actually the smell of chlorine, which emanates from any hypochlorite solution. There is no need to be paranoid about the few hundred mL of gas you have generated, but neither should you be complacent. This element, liberated from ordinary salt, deserves your profound respect. Insert a wet strip of colored paper into your bottle of chlorine and you will see that it immediately bleaches the dye. When you are finished, release your chlorine either outdoors or into a fume hood.

The solution in your chloralkali cell consists of un-reacted sodium chloride, sodium hydroxide, and sodium hypochlorite. The mercury cathode of the Castner-Kellner cell was designed to produce sodium hydroxide in a container separate from the chlorine and salt, preventing the formation of sodium hypochlorite. Without such an arrangement you will be unable to test the pH of your solution using pH test paper because the sodium hypochlorite will bleach the blue color which would otherwise indicate alkali. Instead of testing the pH, you can test the bleaching power of your solution by dipping a strip of colored construction paper into it.

Quality Assurance

In order to pass, your hydrogen must be flammable and your chlorine must smell like chlorine. This is the strong power of all powers, for it overcomes everything fine and penetrates everything solid.

Notes

[1] Goop™ is an excellent choice. There are also epoxies designed specifically for use on plastics.

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Chapter 26. Bayer (Pharmaceuticals)

Alkaloids. There is a considerable number of nitrogeneous compounds, found in plants, which combine with acids to form crystalline salts in the same way that the amines do. Some of these are comparatively simple amines, but most of them are complex in their structure. Many of them have some very marked physiological action as poisons or as medicines. Many which are poisonous are used as medicines in small doses.

Nicotine, C10H14N2, is a colorless oil found in tobacco, which contains from 2 to 8 per cent of the alkaloid. It is very poisonous.

Coniine, C8H17N, the alkaloid of hemlock, is also a liquid. It is historically interesting as the active principle of the fatal draught taken by Socrates.

Atropine, C17H23O3N, is found in Atropa belladonna. It is used to dilate the pupil of the eye and is an active poison.

Cocaine, C17H21O4N, is found in coca leaves. It is used to produce local anaesthesia. A careful study of cocaine has shown that it is a derivative of benzoic acid and the group derived from that acid is chiefly effective in giving to it its valuable qualities. On the basis of this discovery other alkaloids having, in part, a similar structure have been prepared. Some of these retain the anaesthetic effect of cocaine and are less poisonous.

Morphine, C17H19O3N.H2O, is the most important alkaloid of opium and is the chief constituent which gives to laudanum and paragoric their poisonous and sedative qualities. Paragoric also contains camphor and aromatic oils which may have as much effect as the morphine. Opium is obtained from the poppy.

Quinine, C20H24O2N2, is obtained from Peruvian bark. It is a specific in malarial fevers.

Strychnine, C21H22O2N2, is found in Strychnos nux vomica. It is a violent poison, producing convulsions. A dose of 0.06 gram is considered fatal. In small doses it is a powerful stimulant.

— William Noyes, Textbook of Chemistry, 1913 AD [1]

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26.1. Drugs! Now that's a business. Eye of newt, wing of bat, flower of poppy, these are the

stuff that dreams are made of. Up until the twentieth century, pretty much all drugs were 100% natural, meaning they were extracted from plants or animals; just grind up your plant or animal and extract it with hot water or alcohol just the same as if you were making tea or dye. These days, folks tend to think that "natural" is the same thing as "safe." Just try telling that to Socrates or Cleopatra or Napoleon or anyone else who was poisoned before 1900. No sir, some of the most powerful drugs and poisons around are completely natural. But because the plants and animals they come from don't grow just anywhere, drugs that come from far away places tend to be expensive. For example, quinine comes from the bark of the Peruvian cinchona tree. Quinine acts as an antipyretic, reducing fevers in victims of malaria and other tropical diseases. But the cinchona tree won't grow in England, which is why William Perkin tried to make quinine from coal tar in 1856, on account of England had plenty of coal. Of course, he didn't succeed in making quinine; he made mauve instead and that started the whole synthetic dye industry going. But you knew all that from Chapter 22.

The beginnings of the scientific revolution in medicine predate the synthetic dye industry by half a century. Davy had begun his career in 1808 studying the anaesthetic effects of nitrous oxide (laughing gas). Faraday had revealed the anaesthetic properties of ether in 1818. Liebig had prepared chloroform in 1831 and Robert Liston had first used it as a surgical anaesthetic in 1846. Joseph Lister had introduced phenol as a surgical antiseptic in 1867—

Which brings us back to coal tar derivatives. In 1853 Hermann Kolbe synthesized salicylic acid from phenol and carbon dioxide. Salicylic acid turned out to have antipyretic properties similar to salicin, an extract of the bark of the white willow tree. Though cheaper than quinine, salicylic acid had some whopping side effects: nausea, ringing in the ears, and whatnot. Then in 1886 two interns Kahn and Hepp were testing naphthalene for the treatment of intestinal parasites. Patients didn't get any relief from their worms, but the naphthalene reduced their fevers. In a second round of testing, naphthalene killed the worms but didn't help with the fevers. On closer examination Kahn and Hepp found that the first bottle of "naphthalene" had actually been mis-labeled. Consulting the manufacturer, the dye works of Kalle & Company, they found it to be a common dye intermediate, acetanilide. Acetanilide was such a common intermediate, in fact, that there was no hope of patenting it. So instead of marketing acetanilide under its chemical name, as was the custom, Kalle & Company took out a trademark on the brand name, Antifebrin, and began promoting it as an antipyretic. The brand name drug had been born.

Pharmacists, of course, were outraged because they were used to providing generic drugs by their chemical names. But if doctors wrote prescriptions for Antifebrin,

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pharmacists were forbidden by law to substitute what they knew to be the identical and less expensive acetanilide. And no competitive dye house had an incentive to market acetanilide since it could not be protected by patent. The most they could hope for was that one of their own dye intermediates might prove pharmacologically active.

When Antifebrin was introduced by Kalle, Carl Duisberg was the head of research and patenting for Farbenfabriken vormals Friedrich Bayer, one of the many German dye companies. Faced with the disposal of thirty thousand kilos of waste para-nitrophenol, Duisberg challenged his chemists to turn it into a drug. The result was acetophenetidine, marketed by Bayer as an antipyretic and analgesic (pain reliever) under the brand name, Phenacetin. With fewer side effects than Antifebrin, Phenacetin was a cash cow for Bayer, the first of many; in 1898 Bayer introduced a new-and-improved, non-addictive alternative to morphine, then used to treat tuberculosis. Because it made patients feel heroic, Bayer marketed the wonder drug under the brand name, Heroin.

Don't tell me you're going to tell them how to make Heroin! Have you no shame? Have you no sense? Have you no fear of prosecution?

Figment! Did you have a pleasant sulk? I was wondering whether you would ever come back.

You're going to get the Author arrested. What kind of stories do you imagine he'll write for us from the Powhatan Correctional Facility?

Keep your pants on. Heroin is too hard for a beginner to make. I think acetanilide would be a better choice.

I think you ought to have your eight-eyed head examined. Regular use of acetanilide kills red blood cells and turns peoples' skins blue.

Obviously a safer alternative to acetanilide was needed. The year following its introduction of safe, effective Heroin, Bayer came out with a new antipyretic with many fewer side effects than either acetanilide or acetophenetidine. Like its predecessors, acetylsalicylic acid (ASA) had been invented years before and was not patentable under German law. Consequently Bayer chose to market it under the brand name, Aspirin.

The American patent system was a wee bit more accommodating than it was in Germany and the United States had already become the largest market for Bayer dyes. Not only did Bayer have the trademark, Aspirin, it held a US patent. With patent protection, Bayer could sell Aspirin at a higher price than in Europe, and American firms were unable to manufacture the generic ASA. But Bayer recognized that when its seventeen-year patent expired, duties on imported drugs would give American rivals the upper hand. In 1903 Farbenfabriken Bayer began manufacturing Aspirin in New York State through its American subsidiary, the Bayer Company. By manufacturing Aspirin in the United States, Bayer would be able to circumvent the import duties. And by

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promoting Aspirin (not acetylsalicylic acid), Bayer hoped that the brand name would be so firmly entrenched in the medical community that when the patent expired, rival manufacturers of ASA would be unable to compete in the American market. By 1909, Aspirin accounted for 31% of Bayer's US sales.[2]

It was a good plan except for one little glitch; "ethical drugs," the equivalent of modern prescription drugs, were sold through pharmacies, which provided their own packaging. Thus consumers would be unaware of the Bayer trademark. If you wanted to provide your own packaging, you had to put out patent medicines, which were looked down upon by doctors and pharmacists. Ethical drugs had to be tested and proven safe; patent medicines could make whatever claims they wanted and it was up to the customer to decide whether they worked or not. Doctors and pharmacists would boycott any company that tried to sell both ethical drugs and patent medicines. Bayer had positioned itself in the ethical drug market and as such was prohibited from advertising to the general public.

To get around this restriction without offending doctors and pharmacists, Bayer started stamping the Bayer logo into each tablet in 1914. That way, the public would come to know Bayer as the genuine producer of Aspirin. It was pure genius. And it worked. In fact, it worked too well for Bayer's own good. The United States entered the First World War in 1917 and German-owned companies were placed in trust, to be returned at the end of the war. The Bayer Company was cut off from Farbenfabriken Bayer, though its German-born managers were loyal to the parent company. But financial and political scandals resulted in the arrest of the American management team and the seizing of Bayer's American assets. These assets, including the American patents and trademarks, were sold to the highest bidder at the end of 1918, just as the war was ending. The buyer was Sterling Products, manufacturer of such patent medicines as Neuralgine (a headache remedy), No-To-Bac (a nicotine cure), Danderine (a dandruff remedy), and California Fig Syrup (a laxative). Bayer Aspirin was about to go over-the-counter.

In 1919 the co-founder of Sterling, William Weiss, had the temerity to approach Duisberg at Farbenfabriken Bayer with a request for assistance in getting the Bayer Company plants operational, as no one at Sterling knew how to run the equipment. Duisberg had assumed that after the war, Farbenfabriken Bayer would resume its operations in the United States. Now he found that not only had Bayer lost its chemical plants in New York, it had also lost its US rights to the "Bayer" and "Aspirin" trademarks. Even if Farbenfabriken Bayer built new manufacturing facilities in the US, it would be unable to market its ASA as Bayer Aspirin. During the war Duisberg had forged a cartel of the German chemical giants, Bayer, Hoechst, BASF, and AGFA to form IG Farben, of which he was chairman of the board. And now this brash American hay-seed had the audacity to offer Bayer the "opportunity" to sell Bayer products exclusively through Sterling when it became clear that Sterling would be unable to manufacture these products itself.

Now hold on there, cowboy. Sterling had split off all but the Aspirin business into its subsidiary, the Winthrop Chemical Company. Sterling knew everything it needed to

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make Bayer Aspirin. The Bayer Aspirin business was booming, thank you very much, on account of Sterling's advertising genius. It just needed a little help with the sixty-three other Bayer pharmaceuticals to which Winthrop now held the US patents. If Farbenfabriken Bayer would help Winthrop out in the technical department, Sterling would make Bayer a household name.

Bayer was already a household name! Bayer had discovered the analgesic properties of ASA. Bayer had made Bayer Aspirin the universal choice for pain relief. Bayer still held the trademark throughout most of the world. Bayer's American trademarks and patents had been stolen from it during the war; Sterling had simply bought stolen goods. Bayer was happy to reimburse Sterling for its expense, perhaps make Sterling a subsidiary of Bayer. Otherwise, Bayer would see to it that Sterling was known throughout the world as the fake Bayer.

The fake Bayer in Germany, maybe, but in 1919 Sterling bought the rights to Bayer patents and trademarks in England and its sales in South America were booming, thanks to its advertising prowess. No sir, Bayer would deal with Sterling or be relegated to just another producer of ASA. Sterling offered 50% of the Winthrop profits in exchange for technical help. Sterling would have exclusive rights to sell Bayer products, including Bayer Aspirin, in the United States, Canada, Great Britain, Australia, and South Africa. After all, Bayer was good at making; Sterling was good at selling. It would be a win/win situation.

Farbenfabriken Bayer would control the Bayer trademarks and patents over the rest of the world. Sterling would refrain from using Bayer trademarks on Danderine, Fig Syrup, or any of its other snake-oil remedies and would not use Sterling trademarks on any Bayer products. The deal was struck on April 9, 1923 and Bayer Aspirin lived a double life for the next seventy-one years. Then on November 2, 1994 SmithKline Beechambought Sterling Winthrop from Eastman Kodak. The next day it sold the North American over-the-counter component of the business, including the Bayer trademarks, for 1 billion dollars to Miles, Inc., a wholly owned subsidiary of Bayer.

Well, Sterling may have been carved up, scattered to the four winds like so many spiders, but along the way it changed the way that drugs are marketed. Over-the-counter medications are no longer the untested patent medicines of yesteryear. No, these days the market for untested cures belongs to herbal remedies and food supplements. The big drug companies make both prescription and over-the-counter drugs, advertising both kinds of medicine to the general public. And Bayer would begin the twenty-first century as the fourth largest chemical company in the world.[3] Not bad for a dye company.

Notes

[1] Reference [21], pp. 342-343.[2] Reference [64], p. 35.[3] Reference [54], July 29, 2002, p. 16.

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26.2. We have met several classes of organic compounds in previous chapters. Chapter 4 introduced ethanol as the primal member of the alcohols. Alcohols can be distinguished by the presence of an "-OH" functional group, 2/3 of a water molecule, hanging off of various and assorted carbon atoms. Subsequently, Chapter 19 gave us glycerol, a tri-alcohol. As you look on ingredient labels around your house, you can recognized the names of alcohols because they always end in "-ol;" you might find "biglongnamanol," for example.

In Chapter 16 we saw our first organic acid, acetic acid, the compound which makes vinegar sour. Subsequently, Chapter 19 introduced the fatty acids, oleic acid, palmitic acid, and stearic acid. Organic acids are characterized by the peculiar "-COOH" functional group and their names end in "-ic acid;" you might find "biglongnamic acid" or its sodium salt, "sodium biglongnamate" as ingredients in shampoos.

An important concept from Chapter 18 was the acid anhydride. Carbon dioxide, for example, is the anhydride of carbonic acid:

H2CO3 = CO2 + H2O

Sulfur trioxide is the anhydride of sulfuric acid:

H2SO4 = SO3 + H2O

The same idea gives us acetic anhydride, the anhydride of acetic acid:

2 CH3COOH = (CH3CO)2O + H2O

You can think of acetic anhydride as a kind of "super vinegar." It smells like vinegar, tastes like vinegar, and turns into vinegar when you add it to water.

Chapter 22 started us off on a new class of compounds, the amines, of which aniline and toluidine are examples. Just like an alcohol has an "-OH" hanging off of it, an amine has an "-NH2" hanging off of it. You can tell that something is an amine because its name ends in "-ine" or "-in;" biglongnamine would be an example. In fact, if you look back to the beginning of this chapter you will find a whole bunch of natural drugs and poisons—nicotine, morphine, cocaine, quinine—all of them amines. Like ammonia, the amines are bases, or alkalis, and participate in the same kinds of reactions as inorganic bases.

Chapter 7 introduced metathesis reactions, in which inorganic chemicals swapped first and last names. One kind of metathesis reaction was the acid-base reaction, in which a base and an acid combine to produce a salt. Ammonia, for example, reacts with

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hydrochloric acid to form ammonium chloride. The amines can also form salts with acids; while aniline and the toluidines are insoluble in water their salts, aniline hydrochloride and toluidine hydrochloride, are soluble in water. This is a common trick for making pharmaceuticals soluble in water; look at the ingredients of cold remedies and you may find things that look like "biglongnamine hydrochloride," the hydrochloride salt of the otherwise insoluble biglongnamine.

Equation 26-1. From Aniline to Acetanilide

The reaction of an amine with an acid produces a salt, but the reaction of an amine with an acid anhydride produces an amide. Such a reaction, one in which two molecules join together by spitting out a smaller molecule, is called a condensation reaction. Equation 26-1, for example, shows the reaction of aniline with acetic anhydride to produce acetic acid and acetanilide, the drug trademarked as Antifebrin. The dotted lines show where the new bonds will form when acetic anhydride and aniline react. You can imagine half of the acetic anhydride sticking to the nitrogen atom on the amine, while the other half takes off with one of the hydrogen atoms. The acetic acid product is, of course, quite soluble in water, while the acetanilide product is not very soluble and falls out of solution as a precipitate. Thus the acetanilide can be separated from the acetic acid by filtration.

Acetanilide has passed into pharmaceutical oblivion because of its side effects but its near cousin, acetaminophen, has survived as the analgesic trademarked, Tylenol. As a matter of fact, acetanilide is converted into acetaminophen in the body and this latter molecule is responsible for the analgesic effects of the former. We can get the analgesic benefit while avoiding the side effects of acetanilide by taking acetaminophen directly. The structure of acetaminophen is the same as that of acetanilide except that the hydrogen atom marked "-H*" in Equation 26-1 is replaced by an "-OH." In fact, it's this group, characteristic of alcohols, which puts the ol in Tylenol.

The chapter would not be complete without a discussion of the big daddy of analgesics, acetylsalicylic acid, trademarked, Aspirin. Figure 26-1 gives a process schematic for the synthesis of Aspirin beginning with the coal-tar distillate, phenol, an aromatic alcohol.

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You may be familiar with the aroma of phenol, the active ingredient in Chloraseptic throat spray. Reaction of phenol, structure (a), with sodium hydroxide produces sodium phenolate, structure (b). Reaction of sodium phenolate with carbon dioxide produces sodium salicylate, structure (c), and reaction of sodium salicylate with sulfuric acid produces salicylic acid, structure (d). Notice that salicylic acid is a curious beast: it is both an aromatic acid and an aromatic alcohol. Sodium sulfate stays in solution while salicylic acid precipitates and can be filtered out.

Just as the condensation of an amine with an acid anhydride produces an amide, the condensation of an alcohol with an acid anhydride produces an ester. In this case, reaction of salicylic acid with acetic anhydride in toluene solution produces acetylsalicylic acid, structure (e). Acetic acid is soluble in toluene while acetylsalicylic acid precipitates and can be filtered out. Most of the reactants in the Aspirin synthesis, sodium hydroxide, carbon dioxide, and sulfuric acid, are familiar from previous chapters. Similarly, most of the processes are variations on familiar themes, the reaction of an acid with a base and the absorption of carbon dioxide. Only the final reaction, the condensation, is entirely new. This reaction is common to the syntheses of both Aspirin and Antifebrin and is the one we'll explore in this project.

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Figure 26-1. The Aspirin Process

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Material Safety

Locate MSDS's for aniline (CAS 62-53-3), acetic anhydride (CAS 108-24-7), and acetanilide (CAS 103-84-4). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[1]

Your most likely exposure is inhalation of acetic anhydride vapor; if your eyes begin to water, you should go someplace to get fresh air. You should also be aware that aniline may be absorbed through the skin. In case of contact wash the affected area immediately with soap and water.

You should wear safety glasses and rubber gloves while working on this project. All activities should be performed in a fume hood or with adequate ventilation. Leftover aniline should be dissolved in vinegar and flushed down the drain with plenty of water. Leftover acetic anhydride may be dissolved in water and flushed down the drain. Your acetanilide product may be thrown in the trash when you're finished with it.

Research and Development

Before you get started, you should know this stuff. You better know all the words that are important enough to be

indexified and glossarated. You should know all of the Research and Development stuff from

Chapter 21 and Chapter 22. You should know the structures of the coal-tar distillates, aniline and

phenol. You should know the difference between aniline acetate and

acetanilide.

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You should know the structure of acetanilide, acetylsalicylic acid, and acetaminophen.

You should know the hazardous properties of aniline, acetic anhydride, and acetanilide.

You should know the stories of how acetanilide became Antifebrin and how Aspirin went over-the-counter.

Notes

[1] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

26.3. I've made an effort throughout this book to make things out of chemicals that are commonly available, but for this project we'll need aniline and acetic anhydride, neither of which are available "over-the-counter," as it were. Aniline, first encountered in Chapter 22, is a brown oily liquid prepared from coal tar or, more recently, from petroleum. Acetic anhydride is a colorless liquid that smells like vinegar on steroids. Both of these chemicals must be obtained from chemical suppliers as they have no household uses.

In order to make acetanilide, you'll need to answer the usual stoichiometric questions:

Q: Given Equation 26-1 , how many grams of acetic anhydride are needed to react with 1 gram of aniline? We'll need to use excess acetic anhydride because some of it will be hydrolyzed by the water used as a solvent for the reaction. So multiply your calculated weight of acetic anhydride by 1.1 to provide a 10% excess.Q: How many grams of acetanilide should be expected from the complete reaction of 1.0 gram of aniline with excess acetic anhydride? This is the theoretical yield.

Q: Given Equation 26-1, how many grams of acetic anhydride are needed to react with 1 gram of aniline? We'll need to use excess acetic anhydride because some of it will be hydrolyzed by the water used as a solvent for the reaction. So multiply your calculated weight of acetic anhydride by 1.1 to provide a 10% excess.

Q: How many grams of acetanilide should be expected from the complete reaction of 1.0 gram of aniline with excess acetic anhydride? This is the theoretical yield.

Weigh out 1.0 gram of aniline and add it to 10 mL of water in a 100-250 mL beaker. The aniline will sink to the bottom of the beaker, since it's not very soluble in water. Now weigh out your acetic anhydride and add it to the beaker a few drops at a time, stirring the contents of the beaker between drops. In just a few minutes the acetanilide product will

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crystallize. This crude acetanilide is contaminated with acetic acid and aniline acetate. Such contamination cannot be tolerated in a pharmaceutical and it must be removed. Fortunately for us, the solubility of acetanilide differs substantially from that of acetic acid and aniline acetate.

Acetanilide is quite soluble in boiling water, but not in cold water. By contrast, acetic acid and aniline acetate are just as soluble in cold water as in hot water. If you have paid any attention to the previous chapters, it ought to be obvious to you that we need to do some recrystallization. Add another 25 mL of water to the crude acetanilide in your beaker and heat it on a hot plate or over a spirit lamp until the water comes to a boil. All of the solids should dissolve in the boiling water; if they don't, add a little bit more water until they do. Remove the beaker from the heat and set on the counter top to cool.

If you've been earthified to any extent, you'll remember that the slower the recrystallization goes, the purer the crystals will be. So I'm going to suggest that you not rush it; let the hot solution cool slowly and your crystals will be large and pure. The acetic acid and aniline acetate will remain in solution. When the solution reaches room temperature, fold a piece of filter paper, place it into a funnel, and filter your purified acetanilide product. Rinse your crystals with a few mL of cold water, return them to the beaker, and allow them to dry overnight. The spent solution can be poured down the drain. Weigh your dried crystals and express the actual yield as a percentage of the theoretical yield.

Figure 26-2. Acetanilide Crystals

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Quality Assurance

Your acetanilide should take the form of flat, virtually odorless crystals. Record the amounts of materials used and your percent yield in your notebook. Tape a photograph of your crystals into your notebook as a record of your achievement.

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Chapter 27. Badische (Fertilizers)

My chief subject is of interest to the whole world—to every race—to every human being. It is of urgent importance to-day, and it is a life and death question for generations to come. I mean the question of food supply. Many of my statements you may think are of the alarmist order; certainly they are depressing, but they are founded on stubborn facts. They show that England and all civilised nations stand in deadly peril of not having enough to eat. As mouths multiply, food resources dwindle. Land is a limited quantity, and the land that will grow wheat is absolutely dependent on difficult and capricious natural phenomena. I am constrained to show that our wheat-producing soil is totally unequal to the strain put upon it. After wearying you with a survey of the universal dearth to be expected, I hope to point a way out of the colossal dilemma. It is the chemist who must come to the rescue of the threatened communities. It is through the laboratory that starvation may ultimately be turned into plenty.

— William Crookes, The Wheat Problem, 1898 AD [1]

27.1. I don't believe it! A chapter that actually has something to do with air. I suppose the

Author will give it to Samson, who will probably find some roundabout way to work solubility into the discussion. Or how about Lucifer? A little fire and brimstone? No, no, let me guess. The Author will probably assign the chapter to Unktomi, who will undoubtedly earthify the place.

Well, now that you mention it, Crooke's little introductory bit does mention land and soil and such. See, plants need a bunch of elements to grow, hydrogen and oxygen and carbon being at the top of the list. Plants have no trouble getting these things; hydrogen and oxygen come from water and carbon comes from carbon dioxide in the air. From hydrogen and oxygen and carbon they make sugar and from sugar they make cellulose. And what with plants being mostly cellulose and cellulose coming from air and water, I can't blame you for not seeing where the earthification would come in. But besides cellulose, plants need to make proteins and to do that they need nitrogen. Now, you might think plants would have no trouble getting nitrogen, what with the air being 80% nitrogen and all. But most plants can't take nitrogen from the air; they have to get it from the soil. So folks have been using nitrogen-rich dung as a fertilizer since God was a child and it seems to me that earthification is exactly what this chapter needs.

I knew—

Of course, to get from the soil into the roots of the plant, nutrients must be soluble in water. There are two water-soluble nitrogen ions: ammonium and nitrate. Ammonia was

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traditionally supplied by stale urine. Nitrates were traditionally extracted from manure and solublized with potash. So you see, without these soluble nutrients plants would be up a creek.

I knew it was—

Being mere creatures of water and earth and air you undoubtedly failed to appreciate the lessons of Tu-Chhun. The precious nitrates of which you speak are not merely nutrients for vegetables; they feed the arsenals which protect us all from invasion and insurrection. Any society which fails to safeguard its nitrates will inevitably succumb to one which does. Crookes was well aware of this in 1898 and his call for nitrogen fixation was couched in terms of starvation and plenty to make it palatable to those incapable of keeping in the heat and withstanding it.

I knew it was too good to be true. You three are completely missing the point. There you are, surrounded by a virtually unlimited supply of atmospheric nitrogen, but because you don't know how to "fix" it, to convert nitrogen into ammonia or nitrates, you're left to wallow in the dung-heap.

Saltpeter (potassium nitrate) had been the oxidizer of choice for explosives since the invention of gunpowder. In 1846, however, two new explosives burst on the scene; nitrocellulose from the reaction of cotton with sulfuric and nitric acids was discovered by a true Athanor, Christian Sch?in; nitroglycerin from the reaction of glycerine with sulfuric and nitric acids was discovered by another Athanor, Ascanio Sobrero.Nitrocellulose became the basis of smokeless gunpowder, as all of its combustion products are gases. Nitroglycerin remained a laboratory curiosity because of its spectacular sensitivity to shock until a third Athanor, Alfred Nobel, learned to stabilize nitroglycerin with silica. We patented this dynamite, in 1866 and from that time, nitric acid became strategically important as an oxidizer.

Of course, nitric acid was made from saltpeter and sulfuric acid, like it shows in Equation 18-4. Sulfuric acid continued to be manufactured by the lead chamber process, but beginning in 1888 Badische Anilin und Soda Fabrik (BASF) began making it by the more efficient contact process.[2] So acid was as big a business as ever, and saltpeter too, extracted in the traditional way from dung. But from about 1840 folks started importing sodium nitrate mined from huge deposits in Chile, which was exporting over a million tons of the stuff by 1902. But what with wars and famines and population explosions and all, the old-fashioned sources of nitrates weren't going to last forever. Pretty soon everybody and his dog were looking for a way to make nitric acid from something other than saltpeter.

There was ammonia, of course. Available from stale urine in the early days, ammonia became increasingly available in the nineteenth century as a by-product of the manufacture of coke and coal gas. Ammonia could be reacted with sulfuric acid to produce a dandy water-soluble fertilizer, ammonium sulftate. In 1902 Wilhelm Ostwald perfected the oxidation of ammonia over a platinum catalyst to produce nitric acid. From

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ammonia and nitric acid you get ammonium nitrate, the most important fertilizer of the twentieth century, which would usher in an era of unprecedented—

Hold on there, cowboy, you're getting ahead of yourself. Crookes knew all about dung and urine and coal gas and after adding them all up he figured they weren't going to be enough for the industrialized world, much less for the backwaters. Funny story; a fellow named Thomas Willson was trying to smelt calcium from lime long about 1892. It didn't work. Instead of the metal, calcium, he got something that looked like dirt but which turned out to be calcium carbide. When he threw it in the river, up from the water came a bubblin' gas. Acetylene, that is. Acetylene became all the rage for welding and such, which gave Union Carbide its start in 1898.

Fascinating to a hillbilly, I'm sure, but we're talking about fertilizer.

I was just coming to that. So in 1898 Adolph Frank and Nikodem Caro reacted calcium carbide with atmospheric nitrogen in a high-temperature furnace to produce calcium cyanamide, which turned out to be a nifty fertilizer. Pretty soon everybody and her dog were making cyanamid by the Frank-Caro process. American Cyanamid got started that way long about 1907. World production of cyanamid grew from about 7000 tons in 1907 to 120,000 tons in 1913.[3]

The route from cyanamid to nitric acid was not straightforward, however. In 1909 Fritz Haber of BASF reacted hydrogen with atmospheric nitrogen at high temperature and pressure over an osmium catalyst to produce ammonia. The laboratory process was scaled up by BASF engineer Carl Bosch and by 1912 a pilot plant was producing a ton of ammonia per day.

…and just in the nick of time, for Germany. In 1913 Germany imported approximately half of its nitrogen from Chile, a source which became unavailable at the outbreak of WWI due to a successful British blockade. As a consequence, German agricultural and military demands for fixed nitrogen were in direct competition. At the beginning of the war domestic nitrogen came predominantly from the coking of coal; by the end of the war half of German nitrogen production came from the Haber-Bosch process and Germany was nearly self-sufficient in nitrogen. With superiority in fixed nitrogen and aniline dyes, BASF came to dominate German chemicals by the end of the war, and German chemicals dominated the—

Yeah, yeah; almost makes you forget that Germany lost the war, doesn't it? Imported Chile saltpeter, ammonium sulfate from coal gas, and cyanamid were good enough for poor old Britain, France, and the United States. Just imagine if they'd had the Haber-Bosch process. Oh, yeah, I forgot; they did get it—after the war!

In other words, even in the face of German defeat the Haber-Bosch I-dea conquered its competitors during the general economic slump which followed the war. For example, while American Cyanamid assets fell by 8% from 1918 to 1921, BASF assets

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quadrupled.[5] BASF began the twenty-first century as the third largest chemical company in the world.[6]

And Haber was awarded the Nobel Prize in 1919. Twenty-one tumultuous years after Crookes' challenge to the chemical community, virtually unlimited quantities of fertilizer could be manufactured from thin air.

Fertilizers, yes; but it is a mistake to forget that the same chemicals that improve agricultural efficiency so dramatically can have other dramatic effects. The truck bomb that leveled the Murrah Federal Building, for example, was probably filled with ammonium nitrate fertilizer and fuel oil.

I don't believe it; we finally agree about something. At least I don't have to worry about you revealing the secret of high explosives.

Quite the contrary, it is not in the interest of a free society for that particular secret to be kept. An ignorant population is the most vulnerable to attack; in the land of the blind, the one-eyed chemist is king.

Notes

[1] Reference [9], pp. 6-7.[2] Reference [52], p. 85.[3] Reference [61], p. 102.[4] Reference [61], p. 203.[5] Reference [61], p. 254.[6] Reference [54], July 29, 2002, p. 16.

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27.2. Two processes enabled the twentieth century to be one of both unprecedented violence and unprecedented prosperity. The Haber-Bosch process for the synthesis of ammonia made it possible to produce nitrogen fertilizers from air. The Ostwald process made it possible to produce nitric acid from ammonia. Together, these processes freed both farmers and soldiers from their historical dependence on saltpeter, either extracted from animal manure or imported from South America. In the twenty-first century, the production of food and munitions remains subject to political and economic, but not chemical limitations.

Figure 27-1. The Haber-Bosch and Ostwald Processes

The schematic for these processes is deceptively simple. The starting materials, oxygen, hydrogen, nitrogen, and water enter from the left of Figure 27-1, as usual. Oxygen and nitrogen come from the air. Hydrogen might well be supplied by a chloralkali plant, as discussed in Chapter 25, but is more economically generated from natural gas. Water may be obtained from any convenient source. Some engineering is required to separate and purify the starting materials, but the details have been left out of the process schematic for simplicity. The purified starting materials pass through three reactors in turn, a furnace, a burner, and an absorber.

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The first reactor in the schematic, the Haber-Bosch furnace, reacts nitrogen and hydrogen to produce ammonia. This reaction is so simple that you might well wonder why it wasn't developed in the nineteenth century. The reason is that the reaction requires both high temperature (500?C) and high pressure. These extreme conditions were a challenge to engineers trained on the low-pressure reactors of Solvay soda plants, but even with a suitable high-pressure furnace, the reaction goes only at a snail's pace. To speed up the process, hydrogen and nitrogen are passed over pellets of iron oxide and aluminum which serve as a catalyst, a material which makes a slow reaction occur faster but which is neither produced nor consumed in the reaction.

The automotive catalytic converter is, perhaps, the most familiar example of a catalyst. Hot exhaust gases pass over the catalyst, often an alloy of platinum and rhodium. Nitrogen oxides are converted to nitrogen and carbon monoxide is converted to carbon dioxide, reducing the emission of these pollutants from the automobile tailpipe. Unlike gasoline, which is consumed by the car; unlike exhaust, which is produced by the car; the catalyst is unchanged by the reaction. You don't have to refill the converter with catalyst. It is possible, however, for a catalyst to be poisoned, destroying its catalytic properties. The automotive catalytic converter, for example, is poisoned by tetraethyl lead, formerly used as an anti-knock additive in gasoline. Coincidentally, the automotive platinum/rhodium catalyst is also employed in the second reactor of Figure 27-1.

The ammonia produced by the Haber-Bosch furnace is a commodity in its own right, but a large fraction of the ammonia is used to produce nitric acid via the Ostwald process. Except for reactants and products, the Ostwald process is schematically identical to the Lead Chamber process of Figure 18-1. The diverted ammonia enters a burner where it combines with atmospheric oxygen over a platinum/rhodium or gold/palladium catalyst, producing nitrogen monoxide (NO) and water. It should be noted that while a furnace sits over a burner, whose reactants and products are irrelevant to the process, the Ostwald burner is actually fueled with ammonia and it is the exhaust gases themselves which pass to the third reactor in the schematic, an absorber. Atmospheric oxygen is supplied to the absorber to convert these exhaust gases into nitric acid, HNO3. The dilute nitric acid from the absorber can be distilled to produce concentrated nitric acid, which is approximately 70% acid and 30% water. Concentrated nitric acid, like ammonia, is a commodity in its own right, used primarily in the manufacture of fertilizers and explosives.

Black powder is a mixture of charcoal, sulfur, and potassium nitrate, an oxidant which takes the place of atmospheric oxygen. The charcoal fuel can burn much faster than usual, since it doesn't have to wait for fresh air to replace its exhaust gases. But black powder is still just a mixture, separate particles of fuel and oxidant. An even more powerful explosive results if the fuel and the oxidant are present in the same compound. For example, trinitrotoluene (TNT), nitrocellulose (smokeless powder), and nitroglycerin (dynamite) each incorporate fuel and oxidant in a single molecule. Each is produced by the reaction of its base fuel with nitric acid in the presence of sulfuric acid. Nitroglycerin, for example, is produced by the following reaction:

C3H8O3(l) + 3 HNO3(l) = C3H5(NO3)3(l) + 3 H2O(l)

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When nitroglycerin explodes, it does so without requiring a separate oxidizer:

4 C3H5(NO3)3(l) = 12 CO2(g) + 6 N2(g) + 10 H2O(g) + O2(g)

Similarly, nitrocellulose produces no solid combustion products, hence the term, smokeless powder. If nitric acid's only use were the manufacture of high explosives it would be an important chemical, but an even bigger market for nitric acid exists in the manufacture of fertilizers.

Neither nitric acid nor ammonia can be used directly as a fertilizer; nitric acid is a corrosive liquid and ammonia is an alkaline gas. For use as fertilizers both must be neutralized, as shown in the optional reaction of Figure 27-1. Ammonia is often reacted with sulfuric acid to produce ammonium sulfate fertilizer, but since ammonia and nitric acid are produced at the same plant, it's natural to react the one with the other to produce ammonium nitrate. Both the sulfate and the nitrate have neutral pH and make excellent fertilizers, but ammonium nitrate has the advantage of providing a double dose of nitrogen, one in the ammonium ion, the other in the nitrate ion. In fact, ammonium nitrate is just about ideal as a fertilizer except that it's also a powerful oxidizing agent; when combined with a fuel, either accidentally or intentionally, it makes a powerful explosive. An accident aboard the ship, S. S. Grande Camp, caused its cargo of fertilizer to explode in 1947, destroying much of Texas City. Ammonium nitrate is often combined with fuel oil to make a blasting compound for construction, mining, or terrorism. Is ammonium nitrate good, or bad? As with many of the materials we've encountered in this book, ugly is in the eye of the beholder.

Material Safety

Locate MSDS's for cellulose (CAS 9004-34-6), sodium nitrate (CAS 7631-99-4), sulfuric acid (CAS 7664-93-9), nitric acid (400OxCAS 7697-37-2), nitrocellulose (CAS 9004-70-0), ethanol (CAS 64-17-5), and ether (CAS 60-29-7). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[1]

Your most likely exposure is to sulfuric acid. Wash any affected areas immediately with plenty of running water. If eye contact occurs, flush with water and call an ambulance.

You should wear safety glasses and rubber gloves while working on this project. Spent acid solution may be poured down the drain with plenty of water. Nitrocellulose should be burned in a safe place as soon as it is dry. Storage of nitrocellulose is a fire hazard.

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Research and Development

So there you are, studying for a test, and you wonder what will be on it. Study the meanings of all of the words that are important enough to be

included in the index or glossary. You should know all of the Research and Development points from

Chapter 23 and Chapter 24. You should understand the process schematic shown in Figure 27-1

and know the equations given there. You should know the major industries which depend on ammonia and

nitric acid. You should know the hazardous properties of sulfuric acid, nitric acid,

and nitrocellulose, ethanol, and ether.

You should be able to articulate the concerns raised by Crookes and describe the contributions of Ostwald, Haber, Bosch, and BASF, to the fixation of nitrogen.

Notes

[1] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.

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27.3. To make ammonia from nitrogen and hydrogen we require a vessel which can withstand high pressures. For this, we shall use the container that has served us so faithfully in previous projects, the 2-liter soft-drink bottle.

That's it, then; I am living in the head of a madman.

Are you talking to me?

Sorry, I didn't realize you were going to write that down. Apparently, I am talking to you. It seems like a literary device which can only confuse the reader, but what do I know? I'm just a figment of your twisted imagination, after all. But you can't be serious about using a soft-drink bottle to make ammonia. The Haber-Bosch process requires pressures which would turn a pop bottle into a cloud of plastic confetti. Then again, the temperatures required would have turned it into a puddle of goo long before you could fill it with gas.

What do you suggest?

There's a nifty demonstration of the Ostwald process which can be done on a small scale, in a soft-drink bottle, if you like. You put some concentrated ammonium hydroxide into the bottle, heat a piece of platinum wire until it glows red-hot, then plunge the wire into the bottle. It will catalyze the oxidation of ammonia, continuing to glow for a half an hour.[1]

Platinum wire?

This is a private party, if you don't mind.

Pardon me. I am sure that readers will have no problem laying their hands on some platinum wire. Then they will have everything they need to turn household ammonia into nitric acid.

Well, not exactly. It won't make enough acid to actually collect. But the wire will glow, demonstrating—

—the catalytic oxidation of ammonia. Yes, I got it the first time. Figment, you disappoint me. Your early projects were the most boring in the book because there was nothing to make. I thought that you had come to your senses with soap and photography, but now you want to revert to ethereal projects that make no hands—

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—black with charcoal. Yes, I got it the first time. I suppose you have a better idea.

I was thinking of producing nitroglycerin. The glycerine could be distilled from the soap made in one of your previous chapters. It could be nitrated with nitric acid, one of the main subjects of this very chapter, and it would make for an unforgettable demonstration of the importance of nitrogen fixation.

Nah, nitroglycerin is too unstable for the beginner to make. There's no profit in blowing half your reading public to kingdom come.

Finally, the voice of reason—

I think nitrocellulose would be much safer.

I should have known.

Actually, nitrocellulose has a lot going for it. From it, we can make collodion, the material mentioned in Chapter 23 for making glass photographic negatives. Collodion, of course, is a solution of nitrocellulose in ethanol and ether, still used in medicine as artificial skin. Mix collodion with a little camphor and you have celluloid, the first plastic, which foreshadows the subject of Chapter 28.

I suppose it would be too much to ask for the Author to step into the conversation to bring a little order to this literary diversion.

As much as I would like to make chemicals from air, I'm afraid that the engineering concerns are too demanding for us at this point; I think that nitrocellulose is probably a pretty good compromise.

Why not? You've already got them making booze, rockets, and drugs. I'll just sit here quietly and wait for the men in the nice, white suits to come and take us away.

Let's start again. To make nitrocellulose you need to get some concentrated nitric and sulfuric acids.

Folks are going to have a hard time getting nitric acid. You can't just go down to the local hardware store for that.

Hmm. That is a problem. Sulfuric acid can be bought in the form of industrial drain opener,[2] but there's no household use for nitric acid.

You could have them make it from saltpeter, the old-fashioned, pre-Ostwald way.

Instead of making fertilizer from nitric acid, we make nitric acid from fertilizer. Yes, there is a certain poetry in that. So, to make nitrocellulose you need to get some saltpeter,

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either potassium nitrate or sodium nitrate, and some concentrated sulfuric acid. Saltpeter fertilizers are often in the form of "prills," small balls about a millimeter in diameter; if so, you will need to grind your saltpeter to a powder using a mortar and pestle. Weigh 15 grams of sulfuric acid into a small beaker and 10 grams of ground saltpeter into another container. In a fume hood or other area with adequate ventilation, add the saltpeter to the acid and stir it with a glass rod to form a uniform paste. The paste will become warm and if you sniff it cautiously, you'll detect the acrid aroma of nitric acid. Allow the paste to cool while you weigh your cotton.

Figure 27-2. From Cotton to Guncotton

You'll need about 0.5 grams of cotton. In the US cotton is sold as balls weighing about 0.25 grams each. Select two balls as close to this weight as possible and record their exact weights in your notebook. In the following procedure, try to keep each cotton ball intact. When your saltpeter/acid paste has cooled to room temperature, add both cotton balls to the beaker and, using your glass rod, mash them into the paste until they are both thoroughly and uniformly wet with acid through and through. Allow them to sit in the acid for two minutes. Next we need to wash the acid from the cotton. Using your glass rod, fish the cotton balls out of the acid and drop them into a large beaker of water. Use the glass rod to swirl them around in the water, poking them and prodding them as necessary to cleanse them of acid. Dump the acidic wash water from the large beaker into a sink, retaining the cotton with the glass rod. Refill the beaker with fresh water, swirl the cotton thoroughly, and add a spoonful of baking soda, sodium bicarbonate, to neutralize any remaining acid. Test the water with pH test paper; if it still acidic (red or orange) add another spoonful of baking soda and test again. When the water is neutral, pour it down

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the drain, retaining the cotton once again with the glass rod. Refill the beaker a third time, swirl the cotton to wash out the bicarbonate, pour the third rinse down the drain, and place the cotton on some paper towels. Press as much water from the cotton as possible and separate them from one another. Allow one ball of your "guncotton" to dry overnight and use the other to make collodion.

Figure 27-3. From Guncotton to Collodion

To make collodion, we will dissolve our nitrocellulose in a suitable solvent. The traditional solvent consists of 1 part ethanol to three parts ethyl ether. If you don't have access to ether, you can use acetone or ethyl acetate, available as a paint thinner or nail-polish remover (check the label on the bottle). Whatever your solvent, use 10 mL of it in a Petri dish to dissolve one of your cotton balls. Stir the cotton with a glass rod, poking and prodding it to dissolve as much as possible. When it appears that no more will dissolve, fish the remaining cotton from the solvent and dispose of it. Your collodion solution might be used to make photographic emulsion but for now simply allow the solvent to evaporate overnight.

The next day, both your guncotton and your collodion should be dry. Note the resemblance of the collodion film to modern cellophane, also derived from cellulose. Weigh your guncotton and compare its weight to that of the cotton ball from which it was made. Does it weight more or less? Why? Take your guncotton to a suitable location, one free from flammable materials, and carefully light it with a long match. It should burn very quickly, leaving very little ash.

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Ether, acetone, and ethyl acetate are flammable solvents. Pay attention to possible sources of ignition when you are handling them or allowing them to evaporate.Guncotton is a fire hazard and should be stored only in small quantities for as short a time as possible.

Quality Assurance

Your collodion film should be transparent and flexible. Tape it into your notebook and note the similarity of your film to the tape you are using. Unlike modern cellophane tape, collodion film is quite flammable.

Notes

[1] Reference [45], page 214.[2] Be sure to check the label. Household drain openers are usually sodium hydroxide,

but hardware stores often carry drain opener that consists of 95% sulfuric acid.

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Chapter 28. DuPont (Plastics)

When you were in Cambridge this summer you referred to your interest in polymerization. I have some appreciation of the commercial importance of this subject because rubber, cellulose and its derivatives, resins and gums, and proteins may all be classified as large or polymerized molecules. This is a class of substances about which relatively little is known in terms of structure. None of these substances is very amenable to the classical tools of the organic chemist, and no doubt some of the most important contributions in this field will be made by experts in colloidal chemistry. From the standpoint of organic chemistry one of the first problems is to find out what is the size of these molecules and whether the forces involved in holding together the different units are of the same kind as those which operate in holding the atoms in ethyl alcohol together, or whether some other kind of valence is involved—more or less peculiar to highly polymerized substances.

…

For some time I have been hoping that it might be possible to tackle this problem from the synthetic side. The idea would be to build up some very large molecules by simple and definite reactions in such away that there could be no doubt as to their structures. This idea is no doubt a little fantastic and one might run up against insuperable difficulties. The point is that if it were possible to build up a molecule containing 300 or 400 carbon atoms and having a definitely known structure, one could study its properties and find out to what extent they compare with polymeric substances. The bearing of this isn't restricted to rubber but is common to it and such other materials as cellulose, etc.

— Wallace Carothers, Letter to Hamilton Bradshaw, DuPont Experimental Station, November 9, 1927 AD [1]

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28.1. I just want to say one word. Plastics! There's a great future in plastics. It all started in

1845 when a German chemistry professor named Christian Sch?in treated cotton with nitric and sulfuric acids; the result was nitrocellulose, a material that could be molded into just about any shape a body could want. Folks had been using the word plastic since God was a child to describe clay and other such things that could be pressed into shape and it made sense to use the same word to describe this new moldable stuff. Sch?in wrote to Michael Faraday to describe all the cool stuff he was able to make out of it. Well, ten years later news of this cool stuff inspired Alexander Parkes to commercialize the production of nitrocellulose plastics, producing combs and shirt collars and whatnot. From Britain, nitrocellulose crossed the Atlantic, spider-fashion, landing in John Wesley Hyatt, an American who had invented a new way to make dominoes and dice. Billiard balls were next on his list of things to improve and because ivory was expensive, he figured that nitrocellulose might be just the thing. In 1868 he discovered that if you mixed nitrocellulose with camphor (a wood tar), you got a nifty ivory substitute, which he named celluloid, on account of it came from cellulose. Twelve years later, George Eastman came out with celluloid photographic film and in 1891 the first commercial motion picture was produced by none other than the inventor of the electric light bulb, Thomas Edison.

Edison, Edison, Edison. Everybody remembers Edison as the inventor of the light bulb as if he were the only person to succeed in producing one. Nobody remembers that Joseph Swan patiently worked on the light bulb for thirty years. Nobody remembers that he demonstrated a working carbon filament bulb in December of 1878. Nobody remembers that Edison lost his 1882 patent infringement suit against Swan. Nobody remembers that the two inventors went into business as the Edison and Swan Electric Light Company. Nobody remembers that Swan improved the light bulb filament by extruding a solution of nitrocellulose to produce the first artificial fiber. Nitrocellulose would make its mark as the first artificial silk substitute, but its extreme flammability made it less than ideal as a textile. Continuing the quest for perfect light bulb filament, Edward Weston dissolved cotton in ammoniacal copper hydroxide producing cupprammonium rayon. And by a curious coincidence Charles Cross and Edward Bevanwere seeking the Holy Grail of light bulb filaments when they invented viscose rayon and cellulose acetate. In 1910 Henri and Camille Dreyfus of Celanese Corporation began manufacturing cellulose acetate film for those who did not appreciate the tendency of nitrocellulose movie film to go up in smoke.

I beg your pardon. Eastman's celluloid film may have been flammable, but at least it was smokeless. A century before the invention of acetate film Ir 鮩 e du Pont de Nemours had immigrated to the United States to avoid the aftermath of the French Revolution. Having apprenticed under Lavoisier, du Pont was well-versed in the manufacture of gunpowder and, finding that American gunpowder was both expensive

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and sub-standard, he constructed a gunpowder mill in 1802. Demand for gunpowder increased dramatically throughout the nineteenth century as mining operations marched across the North American continent. The wars of that century were also a boon to powder manufacturers; the du Pont de Nemours Powder Company (DuPont) supplied powder to both sides in the Crimean war and provided one third of the explosives used by the United States during the Civil War. After the war, DuPont began to explore explosives based on nitroglycerin and nitrocellulose. To circumvent Alfred Nobel's dynamite patents, DuPont bought the California Powder Company in 1876. California Powder produced a mixture of nitroglycerin, sugar, and saltpeter called "White Hercules." At the close of the nineteenth century DuPont moved into nitrocellulose-based "smokeless powder" and held a monopoly in its supply to the US military.

Monopoly was DuPont's favorite game, as a matter of fact. By 1881 the Gunpowder Trade Association, or "Powder Trust" controlled 85% of the US gunpowder market and DuPont controlled the trust. Then in 1907 DuPont became the target of an anti-trust suit, the first of many, which discouraged its erstwhile practice of growth through the acquisition of its competition. In 1913 DuPont was broken up into the Hercules Powder Company, the Atlas Powder Company, and the du Pont de Nemours Powder Company. Hercules and Atlas were to compete with DuPont in dynamite and black powder; DuPont managed to hold on as the sole provider of smokeless powder to the US military. With plenty of cash from World War I ammunition sales, DuPont looked for expansion opportunities which would provide a peace-time outlet for nitrocellulose without setting off an anti-trust inquisition. In 1915 DuPont bought the Arlington Company, the largest American producer of celluloid and celluloid products, and two years later it acquired Harrison Brothers, a firm specializing in paints, pigments, and chemicals. DuPont developed a tough nitrocellulose automotive enamel, trademarked "Duco," which General Motors adopted in 1922 as a colorful alternative to Ford's basic black. From nitrocellulose enamel, DuPont moved into other cellulose products, notably viscose rayon, cellophane, and acetate film. In 1924 DuPont began manufacturing synthetic ammonia (from air), ending its dependence on imported nitrates as a source of essential nitric acid. DuPont was fast becoming a diversified chemical giant; by 1933 90% of its earnings would come from chemicals other than explosives.[2]

It was for such a company that I, Wallace Hume Carothers, went to work in 1928. I was the egghead kid everyone made fun of in school—glasses, pocket protector, the whole shebang. My father sent me to study accounting at junior college despite my interest in science. With the depression on, I managed to land a job teaching business at Tarkio College, where I was able to take chemistry classes on the side. I soon found that my happiest times were when I was left alone in the lab. For the unhappy times I was comforted by a small vial of cyanide I kept on my watch chain. It was my little secret. Hoping to make something of myself, I enrolled in the graduate chemistry program at the University of Illinois. After graduation, Harvard gave me a job teaching chemistry, but I was a nervous lecturer and longed to be left alone with my research, my record player, and the odd bottle of beer. I soon accepted a position as head of research at DuPont, hoping that without the distractions of teaching I could finally succeed in doing

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something important with my life. And if I couldn't do something important, there was always my little secret.

Lighten up, man, you're creeping me out. Illinois Ph.D., Harvard professor, DuPont bigwig, what more could you want out of life?

The world was changing and I wanted to make a contribution. Most of the history of chemical industry had involved breaking big molecules down into smaller ones; this is what destructive distillation is all about. But chemists were just beginning to combine small molecules into big ones. In 1906 Leo Baekeland had introduced the plastic Bakelite, a polymer of phenol and formaldehyde. Julius Nieuwland had succeeded in polymerizing acetylene in 1925 and DuPont had licensed the process. My research team followed this up by reacting vinylacetyene with hydrogen chloride to form the monomer, 2-chlorobutadiene. This small molecule polymerizes into an elastic material, trademarked Neoprene by DuPont in 1933. Neoprene was more resistant to petroleum distillates than natural rubber, but the depressed price of natural rubber made neoprene too expensive for practical use.

What about polyester?

What about it? Nobody at the time knew what held polymers together. Some supposed that polymers were simply loose aggregations of smaller molecules; others believed that these smaller molecules were chemically bonded to one another. It seemed to me that if it were possible to build long chains of short molecules using known chemical reactions, and if the properties of these long chains were similar to those of natural polymers, these observations would support the "chemically bonded" hypothesis of polymer structure. It was my one good idea. To this end, my group began studying the condensation of acids and alcohols to produce esters, the same reaction that produces ethyl acetate from ethanol and acetic acid. We began our work on these polyesters by reacting ortho-phthalic acid, a molecule with two acid groups, and ethylene glycol, a molecule with two alcohol groups. We produced some interesting fibers but their melting point was too low to be useful as a textile. We gave up on it in 1930, but thanks for bringing it up.

Stop your whining. The rayons had been merely artificial fibers derived from natural cellulose. Your polyethylene orthophthalate was the first truly synthetic fiber, one whose starting materials were not polymers to begin with.

Since the polyesters had not worked out, we started on polyamides. We tried many acids and amines derived from petrochemicals, but the 6-carbon adipic acid and the 6-carbon hexamethylene diamine seemed to work out the best and DuPont trademarked this polymer as Nylon. Even so, I was unhappy. Perkin had invented mauve before the age of twenty; Hall and H 鲯 ult had invented electrolytic aluminum when they were both twenty-three; Davy had discovered three elements by the time he was thirty-two. At forty-one I had had only one good idea and I feared that I would never have a second one. I spent my days churning out too many corporate reports and my nights knocking down

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too many bottles of beer. As a mortal I had kept my little vial of cyanide a secret, but in 1937 that particular secret would no longer be kept. Hello Darkness, my old friend.

It seems to me you lived your life like a spirit lamp in the wind.

Unbelievable! You had everything to live for.

Quite the contrary, Figment; he had nothing to live for. He had had his one great I-dea and it no longer needed him. Spreading through the scientific literature and the corporate culture, it had taken on a life of its own. So when no hope was left in sight on that starry, starry night, he took his life as watery chemists seem to do.

Well, if he had stuck around he would have seen Nylon become DuPont's single largest source of income in the fifties and sixties. His polyester recipe would be modified slightly to produce polyethylene terephthalate, trademarked by DuPont as Dacron polyester. DuPont would enter the twenty-first century as the Earth's second largest producer of "better things for better living through chemistry."[3]

Notes

[1] Reference [6].[2] Reference [73], p. 143.[3] Reference [54], July 29, 2002, p. 16.

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28.2. We began this book with the combustion of cellulose, a polymer of glucose, and progressed to spinning yarn from protein fibers. We have selectively dissolved the polymer, lignin, from plant stems to produce paper and we have converted cellulose into a different polymer, nitrocellulose, to produce collodion film. If we wished we could dissolve cellulose in a solution of copper sulfate and ammonia to produce cuprammonium rayon,[1] but we would merely be making one polymer from another. We might also try our hand at Bakelite or neoprene, but these polymers were created "in the dark," so to speak, with no real understanding of what it was about these substances that gave them the properties commonly associated with polymers. No, polyesters were the first polymers to be created "on purpose," that is, with some clue as to what structural features of a monomer might promote polymerization. For this reason we shall peruse polyesters in particular, pointing out principles also applicable to polyamides and proteins.

Carothers' greatest contribution to science was to suppose that polymers might be designed using known chemical reactions. One of the kinds of reactions we have seen time and again has been the condensation reaction, one in which two molecules join together, spitting out a small molecule such as water. We saw cellulose condense from glucose, and protein from amino acids in Chapter 6, ethyl acetate from ethanol and acetic acid in Chapter 16, fats from glycerol and fatty acids in Chapter 19, and acetanilide from aniline and acetic acid in Chapter 26. Carothers did not know that cellulose and proteins were condensation polymers, but he did know about the structure of fats, ethyl acetate, and acetanilide. He wondered whether it might be possible to build up large molecules from small ones using condensation as the means for linking them together.[2]

How, then, might it be possible to hook more than two molecules together? When ethanol reacts with acetic acid, the -OH part of the alcohol reacts with the -COOH part of the acid. The product, an ester, has neither an -OH nor a -COOH group with which to hook onto a third molecule. But what if we could devise an alcohol with two -OH groups? Then it would be possible to hook an acid onto each of them, forming a larger molecule from three smaller ones. Similarly, if we could devise an acid with two -COOH groups, it would be possible to hook an alcohol on each of them, again forming a larger molecule from three smaller ones. Now here comes the genius part; what if we were to react a di-alcohol, one with two -OH's, and a di-acid, one with two -COOH's? Well, when the di-alcohol hooked onto the di-acid, the resulting ester would still have an -OH at one end and a -COOH at the other. The -OH could hook onto another di-acid, whose -COOH could hook onto another di-alcohol, whose -OH could hook onto—well, you get the picture. The condensation product of a di-alcohol and a di-acid could be very long indeed. And if these long molecules were to have properties similar to those of known polymers, if they could be drawn into fibers or stretched into films, we might reasonably conclude that they were new polymers, polyesters.

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So to make our polyester, we need a di-alcohol. We have already been introduced to a tri-alcohol, glycerol, in Chapter 19. The beginning of the twentieth century saw increased production of the most common di-alcohol, ethylene glycol, introduced in 1927 by Union Carbide as an automotive anti-freeze.[3] Ethylene glycol began to be manufactured from petroleum rather than coal tar and from the roaring twenties on, petroleum companies like Standard Oil and Shell began to compete with established chemical manufacturers. The anti-freeze used today is essentially pure ethylene glycol with a little dye added to make leaks easier to spot. Ordinary anti-freeze will be our source of ethylene glycol.

We also need a di-acid; from the last quarter of the nineteenth century, one had been produced by the coal-tar industry for the manufacture of dyes and pharmaceuticals. Phthalic acid was traditionally produced by the oxidation of naphthalene, a coal-tar distillate familiar as the stuff from which moth balls are made. As with other chemicals, oil companies learned to make benzene, toluene, phenol, aniline, naphthalene, and other traditional coal-tar products from petroleum and today both ethylene glycol and phthalic acid are manufactured primarily from oil rather than from coal. We will actually find it more convenient to use phthalic anhydride for our condensation, just as we used acetic anhydride for the condensation of acetanilide from aniline in Chapter 26. I suppose that you could try making your own phthalic anhydride from moth balls, just as you made nitrocellulose from cotton in Chapter 27, but otherwise you will have to get it from a chemical supply company.

Equation 28-1. The Condensation of an Ester

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Equation 28-1 shows what happens when ethylene glycol reacts with phthalic acid. The product, ethylene glycol monophthalate, has an -OH at one end and a -COOH at the other, just as we imagined a few paragraphs ago. When two ethylene glycol monophthalates get together, they undergo consensation, spitting out water and joining together to form a larger molecule, a di-ester. A free -OH remains at one end and a free -COOH at the other and the process of polymerization can continue indefinitely to form a polyester, polyethylene phthalate.

The properties of phthalic acid lend themselves to the laboratory preparation of a polyester, but the melting point of this particular polyester is too low for commercial use as a fiber or film. A very small change of reactants, however, produces polyethylene terephthalate, trademarked as Dacron for use as a fiber, and Mylar for use as a film. Polyethylene terephthalate is quite familiar to caveman chemists; it is the stuff from which the 2-liter soft-drink bottle is made. The "very small change" is the substitution of terephthalic acid for phthalic acid. These two compounds are isomers of one another; they have the same molecular formula but different structures. In phthalic acid the -COOH groups are next to each other, in terephthalic acid they are on opposite sides of the aromatic ring. Other than the conditions of temperature and pressure used for the commercial synthesis, the polymerization of polyethylene terephthalate is the same as that of its cousin, polyethylene phthalate.

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Figure 28-1. Dacron, Nylon, Protein, and Cellulose Monomers

Polyesters were the first synthetic fibers to be prepared by Carothers and his team at DuPont, but their first commercial success in this area was the polyamide, Nylon. Whereas polyesters are condensation products of acids and alcohols, polyamides are condensation products of acids and amines. Just as different polyesters result from the choice of acid and alcohol, different Nylons result from the choice of acid and amine. The

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most common Nylon, Nylon-66, results from the condensation of the six-carbon adipic acid and the six-carbon hexanediamine. At least one Nylon, Nylon-6-10, can be conveniently produced in the laboratory.[4]

Two natural polymers, protein and cellulose, have played a prominent role in the history of technology. Figure 28-1 compares the structures of these compounds to their synthetic progeny. Each structure shows two monomer units with the exiting water molecule circled. The proteins are a particularly varied class of polymers because there are twenty amino acid monomers to choose from and they can be strung together in a bewildering number of combinations. The monomer chosen for the figure is the amino acid glycine. Cellulose is a polymer of glucose, one of the principle sugars in honey. It has been central to the story of textiles and paper, industries which created the demand for dyes and alkalis. It gave us collodion, celluloid, and rayon as antecedents to the synthetic polymers. It is fitting that we should end our exploration of chemistry here, for cellulose was the material from which we started the fire which launched us on our long journey from caveman to chemist way back in Chapter 1.

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Material Safety

Locate MSDS's for ethylene glycol (CAS 107-21-1) and phthalic anhydride (CAS 85-44-9). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[5]

Your most likely exposure will be to spilled ethylene glycol. Wash your hands when you have finished this project.

You should wear safety glasses while working on this project. Leftover solid materials can be thrown in the trash. Leftover ethylene glycol should be saved for re-use or collected for disposal wherever automotive products are sold. Ethylene glycol, like all alcohols, is toxic by ingestion. Its sweet taste makes it particularly attractive to unsuspecting animals and so care must be taken in its disposal.

Research and Development

You are probably wondering what will be on the quiz. You should know the meanings of all of the words important enough to

be included in the index or glossary. You should know the Research and Development points from Chapter

23 and Chapter 24. You should know the hazardous properties of phthalic anhydride,

ethylene glycol, and sodium acetate. You should know the reaction of Equation 28-1. You should be able to sketch the structures shown in Figure 28-1. You should be able to describe the condensation of esters and amides.

You should be able to tell the story of Wallace Carothers.

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Notes

[1] Reference [44], page 247.[2] There are a bewildering variety of synthetic polymers, of which the condensation

polymers make one class. The other class, the addition polymers, includes polyethylene, polyvinyl chloride, polystyrene, polyisoprene (natural rubber), neoprene, Teflon, Orlon, and Plexiglas.

[3] Reference [61], p. 373.[4] See Reference [44], p. 213.[5] The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or

Saf-T-Data ratings at your convenience.

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28.3. The method you will use in this project comes right out of the journal article which introduced the word, polyester.[1] It is not a commercial polyester. It melts at too low a temperature to be used as a textile, but its low melting point makes it convenient for us to make in the laboratory without special equipment. To make a polyester you need a di-alcohol and a di-acid. One di-alcohol, ethylene glycol, is readily available from automotive suppliers as anti-freeze. When shopping for anti-freeze, check the label; you want a close to pure ethylene glycol as possible. Look for a generic, no-frills, non-environmentally-friendly, undiluted anti-freeze, the cheaper the better. You do not want pre-mixed anti-freeze because it has been mixed with water and you will be trying to eliminate water from the reaction.

Unfortunately, there is no consumer market for di-acids so you will have to get one from a chemical supply. Phthalic anhydride is the anhydrous form of phthalic acid, which is particularly convenient for condensation reactions since you are trying to eliminate water from the reaction. The reaction is:

n C8H4O3 + n C2H6O2 = (C10H8O4)n + n H2O

Of course, you will need to answer the stoichiometric questions:

Q: How many grams of phthalic anhydride will react with 1.0 g of ethylene glycol?Q: How many grams of polyethylene phthalate should be expected from the reaction of 1.0 g of ethylene glycol?

Q: How many grams of phthalic anhydride will react with 1.0 g of ethylene glycol?

Q: How many grams of polyethylene phthalate should be expected from the reaction of 1.0 g of ethylene glycol?

You are probably wondering what to do with the "n" in the equation. Since the stoichiometry involves only the ratios of products and reactants, the "n" will cancel out of any stoichiometric calculation. Try it with n=1 or n=10 or n=100; you should get the same answers no matter what choice you make for "n."

Before you begin, you will need to know the weight of your empty test tube. It is convenient to place a small beaker on your balance, tare it so that it reads zero, then place the empty tube into the beaker and write down the weight of the empty tube in your notebook. The use of the beaker will allow you to weight the tube later on even when it is full of liquid. Using the methods of Appendix C, weigh your calculated amount of phthalic anhydride and 1.0 g of ethylene glycol into your test tube.

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Light your spirit lamp and use it to heat the test tube. You will, of course, need to hold the test tube with a test tube holder. It would also be wise to wear a leather work glove on your test-tube hand since even the holder may get hot. As you heat the tube you want to be sure to heat it evenly from top to bottom. Hold the tube horizontally as much as is possible without losing its contents. Move the tube back and forth in the flame, turn it over, and heat the other side. Your goal it to heat the entire tube, not just a part of it. As it warms, the phthalic anhydride will begin to melt and dissolve in the ethylene glycol. You may notice that as it does so, the solution appears to boil; the water produced by the reaction is boiling away. If the mouth of the tube is cool, the water may condense and dribble back into the reaction. We heat the entire tube evenly to prevent this from happening. You might think that more heat would be better, but at this stage our goal is to boil away the water, not the ethylene glycol. To this end, you should heat your tube gently and with great skill. If the reaction begins to boil vigorously, remove the tube from the flame for a moment until it settles down.

When all of the solid anhydride has melted, the production of water will noticeably decrease. Continue to heat it gently for another minute or two and then weigh the tube, taring the beaker as before, and record the weight in your notebook. Subtract the empty weight of the tube from the full weight, record the actual yield of polyester and express it as a percentage of the theoretical yield. If this percentage is over 105%, you have not eliminated all of the bejeesical water; return the tube to the flame for another round of gentle heating. If the percentage is below 95% you have boiled away some of your ethylene glycol prematurely; you may continue, but you will have less polyester than you might have had.

When your yield is at or below 100%, return the tube momentarily to the flame until any solid material has melted once again. Then pour the hot polyester into the middle of a square of aluminum foil. Touch your glass rod to the hot, viscous polyester. A little of it will stick to the glass rod and dribble back into the puddle of molten plastic. Touch the glass rod to the polyester and you will notice that it becomes more viscous as it cools. Continue touching the rod to the polyester until you are able to draw a fiber from it, as shown in Figure 28-2. This fiber will be short in the beginning, perhaps a few cm, because the polymer chains are not very long. In other words, the "n" is some small integer. Eventually, the polymer will cool to the point that you can no longer draw a fiber from it.

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Figure 28-2. Drawing a Polyester Fiber

Imagine the molecular dating scene at the beginning of the reaction. As each phthalic anhydride molecule melted, it would find itself surrounded by unattached ethylene glycols. Each one would hook up with two ethylene glycol molecules, since bigamy is the natural state for a phthalate. Toward the middle of the reaction, a melting phthalic anhydride molecule would find that many of the free glycols were already taken. No problem—each of the glycols attached to the earlier phthalates would still have had a free -OH available for grabbing. Occasionally, two short chains would find one another and hook end-to-end to form a longer chain, but toward the end of the reaction the chains would be long enough that the loose ends would almost never find one another. Furthermore two ends would have to be of the opposite "gender" to hook up properly; two glycol ends would be out of luck in the condensation department. Consequently, the longer the chains get, the harder it would be for them to grow.

If you want long fibers, you need long chains. To promote this, we can return the polymer to the flame, this time by holding the aluminum foil over the spirit lamp with gloved hands until the polyester melts once more. Continue heating gently and you will start to boil off any ethylene glycols and phthalic acids which happen to fall off the ends of their respective chains in the heat of the moment. If two glycol ends were near one another initially and one of the glycols were boiled away, the other would now be free to hook up with the now-available phthalate end of the other chain. Allow the polymer to cool and you will find that you can draw longer fibers from the melt than you could before this second heating. Carefully wind your fiber onto the glass rod to form a small cocoon, as

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shown in Figure 28-2. You may return foil to flame as many times as you like but be careful not to scorch the polymer; it is possible to produce fibers 10 meters long by repeated heating of the polyester.

Quality Assurance

Record in your notebook the percent yield and the length of your longest polyethylene phthalate fiber. Sadly, your fibers are likely to be too fragile to keep permanently; photograph your best fiber as a kind of phthalic symbol. This photograph and the polymer remaining on the foil should be taped into your notebook.

Notes

[1] Reference [6].

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Appendix A. Back Cover

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Appendix B. The Laboratory NotebookRe-minding is a tedious process. Starting with a virtual blank slate for a mind, it takes

two years to learn to walk, another five or ten to condition the animal body for effective communication, and another twenty or sixty or even eighty years to assemble the requisite thoughts, notions, values, skills, observations and visions to form a complete human being. And just when you have everything in its place, your animal host gets itself trampled by a herd of water buffalo or killed in battle or it simply wears out and the whole process has to start over again with a fresh animal host.

Students at Hampden-Sydney College are required to keep a notebook documenting their achievements as caveman chemists. The notebook serves the purpose of preparing them for the work, organizing it, and documenting it. If they ever desire to repeat a project in the future, the notebook should provide a sufficient level of detail for them to do so. If a project does not turn out, the notebook helps me to help them to figure out what went wrong. It also serves as a souvenir for the children and grandchildren.

Caveman chemists are encouraged to follow the Hampden-Sydney format:

The inside front cover has a table of contents listing all of the potential projects and the page numbers at which they appear in the notebook. When you are ready for a project to be graded, initial its entry in this table of contents.

You should use only ink in the notebook so that information is not easily lost. Mistakes receive special attention in a notebook of this kind. Your impulse may be to obliterate a mistake, but you may decide later that what you thought was a mistake was correct after all. If you have obliterated your mistake, or torn it out of your notebook, you have no way to recover it. For this reason, simply draw a single line through a mistake. This marks it as a mistake, but allows it to be recovered later if necessary.

Number each right-hand page in its upper right-hand corner. Whenever you begin work on a page, write the date next to the page number.

There is potential confusion between scratch calculations and data actually observed in the laboratory. To avoid such confusion, always record experimental observations on a right-hand page. Reserve left-hand pages for scratch work, estimates, and theoretical predictions.

Begin each project on a new right-hand page. Give each project a title and organize it into the following sections:

1. Purpose: Write a paragraph describing the purpose or importance of the project.

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2. Materials: Make a bulleted list of all the materials needed to complete the project.

3. Procedure: Translate the instructions for each project into a bulleted list which you can check off, item by item, as you work on the project. Some cavemen have suggested that I provide them with such a list, but doing it yourself increases the chances that information will actually pass through the brain on its way from the eye to the hand. The left-hand page facing the Procedure may contain calculations to estimate the amounts of materials needed. Record the results of these calculations along with the rest of your plans. If you are working in a class, you may have been quizzed on the material given in the sections labeled "Research and Development." If so, tape the successful quiz to the left-hand page opposite the Procedure.

4. Safety: A brief summary of the hazardous properties of the materials and equipment to be used, including first-aid for accidental ingestion or exposure to eyes, skin or lungs.

5. Observations: A brief summary of the deviations from the planned Procedure. The actual amounts of materials, as distinct from the estimated amounts, should be recorded here.

6. Results: A brief description of the material or item produced. Record any relevant properties such as weight, color, and structural integrity.

7. Conclusions: A paragraph summarizing what you have learned from the project. Were you successful? How do you know? What might you do differently if you were to repeat the project? How does this project relate to others you have done or may plan to do?

The inside back cover has a list of completed projects initialed by the professor.

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Appendix C. Measuring and MixingOne of the most important skills any chemist must master is the ability to accurately weigh out chemicals without contaminating them. You really need two balances[1] for the projects in this book, one for weighing big things and one for weighing little things. The big balance we'll call a gram balance because it can weigh things to the nearest gram. These jobbies usually have a capacity of 2 kg, which means you can load them down with 2 kg of stuff before they max out. The little balance we'll call a centigram balance; it can weigh to the nearest 0.01 g and usually has a capacity of 200 g. You'll use the centigram balance whenever the thing you are weighing weighs less than 200 g and the gram balance whenever the thing you are weighing weighs between 200 and 2000 g.

You might think weighing things out would be simple; just put a container onto a balance and spoon stuff into it until it reads the right amount. There are two problems with this approach. Let's say that you are spooning silica and plaster into a plastic bag. You spoon in the right amount of silica but you get a little too much plaster. What do you do? You reach in and spoon out a little plaster. But plaster and silica look pretty much the same, so how do you know which one you took out? The second problem is that you might be tempted to put the extra "plaster" back into your container of plaster. After all, why waste it? But chances are you just spooned a little silica back into your container of plaster. Over time all those little bits of contamination will add up and pretty soon all your chemicals will be contaminated. There's a simple method for avoiding these problems. It's called weighing by difference.

Let's go back to the example. You have a container of silica and a container of plaster. Your goal is to get, say, 20 g of each into a bag. First place the container of silica on the balance and press the button marked "tare."[2] The balance will now read zero. Spoon some silica from the container into the bag. The balance will now read, say, -2 g because the container has 2 g less silica than it had when you tared the balance. Suppose you keep spooning and with the last spoonful the balance reads -22 g. What do you do? The spoon is still sitting over the container on the balance. It hasn't been anywhere to get contaminated. Just flick it with your finger until the balance reads -20 g and then dump the last spoonful into the bag. If you flick too much and get -19 g, just reach in and grab a little more until you get it just right. Now clean off your spoon, take the container of silica off the balance, put the container of plaster on the balance, and press the tare button again. It reads zero, of course. Spoon some plaster into your bag. Suppose that with the last spoonful the balance reads -23 g. The spoon still hasn't been anywhere to get contaminated. Just flick it with your finger until it reads -20 g and dump the last spoonful into the bag. Your silica container still contains silica and your plaster container still contains plaster. Your bag contains exactly 20 g of each. Weigh by difference and you will keep all of your chemicals as pure as the day you made them.

It may seem strange to you that I had you weighing things into a bag, but that's a great way to mix dusty materials. With silica and plaster, for example, you can seal the plastic

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bag and then knead it, shake it, massage it, and turn it end for end until the silica and plaster are all mixed up and you won't get dust everywhere. If the powders are flammable, I like to use those anti-static plastic bags like circuit boards come in. That way, you won't have static electricity accidentally set it off.

I buy silica in 50-pound bags and you might wonder how I can weigh it out if the balance has a capacity of only 2 kg. I use 1-pint plastic tubs like they use at Chinese restaurants to store each of the chemicals I use frequently. I label each container and its lid so that the silica container never gets used for anything else. These containers are just the right size for a gram balance. I keep large amounts of chemicals in 5-gallon pails, like paint comes in. I always take chemicals from a big container into a smaller container, from the 5-gallon pail to the 1-pint tub, from the 1-pint tub to the plastic bag. As long as I keep this in mind I will never contaminate more than a little bit of chemical.

You can weigh liquids by difference just as you can solids. Just pour the liquid into a cup, place the cup on the balance, tare it, and use a medicine dropper to transfer liquid from the cup to wherever it's going. But we often measure liquids by volume rather than by weight. You can use a graduated beaker or measuring cup for large volumes, graduated a graduated cylinder for smaller ones, and a graduated pipette for really small volumes.

Notes

[1] What most people call a "scale," chemists call a balance. You may use either mechanical or electronic balances.

[2] On some balances the tare button is marked "zero."

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Appendix D. Supplies and Suppliers

Ammonium Hydroxide (Clear Ammonia) Grocery Store

Acetic Acid (Vinegar) Grocery Store

Acetic Anhydride Chemical Supply

Aniline Chemical Supply

Borax Grocery Store

Calcium Carbonate (Limestone) Garden Supply, Farm Supply, Ceramic Supply

Calcium Hydroxide (Slaked Lime) Garden Supply, Farm Supply, Grocery Store

Calcium Magnesium Carbonate (Dolomite)

Farm Supply, Garden Supply

Calcium Sulfate Dihydrate (Gypsum) Garden Supply, Farm Supply

Calcium Sulfate Hemihydrate (Plaster of Paris)

Ceramic Supply

Carbon (Charcoal) Nature, Grocery Store, Pyrotechnic Supply, Pet Store

Copper Carbonate (Malachite) Ceramic Supply

Copper II Oxide, Black Ceramic Supply

Ethanol (Ethyl Alcohol) Building Supply, Liquor Store

Ethyl Ether Chemical Supply

Ethylene Glycol (Anti-Freeze) Automotive Supply

Indigo Weaving Supply

Phthalic Anhydride Chemical Supply

Potassium Carbonate (Pearl Ash, Potash) Ceramic Supply

Potassium Nitrate (Saltpeter) Garden Supply, Farm Supply, Pharmacy, Pyrotechnic Supply

Silicon Dioxide (Silica, Flint, Sand) Ceramic Supply

Silver Nitrate Photographic Supply, Chemical Supply

Sodium Carbonate (Soda Ash) Ceramic Supply

Sodium Carbonate Decahydrate Grocery Store

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(Washing Soda)

Sodium Hydrosulfite (Color Remover) Grocery Store, Weaving Supply

Sodium Hydroxide (Caustic Soda, Lye, Drain Opener)

Grocery Store, Plumbing Supply

Sodium Hypochlorite (Laundry Bleach) Grocery Store

Sodium Nitrate (Chile Saltpeter) Farm Supply, Garden Supply

Sodium Thiosulfate (Hypo) Photographic Supply, Chemical Supply

Sulfur Garden Supply, Farm Supply, Pharmacy, Pyrotechnic Supply

Sulfuric Acid (Drain Opener) Plumbing Supply

Tin Oxide (Cassiterite ) Ceramic Supply

o-Toluidine Chemical Supply

p-Toluidine Chemical Supply

Wool Weaving Supply

Suppliers

Caveman Supply Ceramic Supply

Campbell Ceramic Supply , 5704 D General Washington Dr., Alexandria, VA 22312, (800) 657-7222

Clay Art Center , 2636 Pioneer Way East, Tacoma, WA 98404, (253) 896-3824

Chemical Supply Antec Inc. , 721 Bergman Avenue, Louisville, KY 40203, (800) 448-2954 Voigt Global Distribution , PO Box 412762, Kansas City, MO,64141-

2762, (816) 471-9500 The Chemistry Store , 520 NE 26th Court, Pompano Beach, FL, (800) 224-

1430 Cynmar Corporation , PO Box 530, Carlinville, IL 62626, (800) 223-3517 Tri-Ess Sciences , 1020 Chestnut Street, Burbank, CA 91506, (800) 274-

6910 The Al-Chymist , 17130 Mesa Street, Hesperia, CA 92345, (760) 948-4150

Homebrew Supply Beer and Wine Hobby , 155 New Boston St., Unit T, Woburn, MA 01801,

(800) 523-5423

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The Home Brewery , 205 West Bain PO Box 730, Ozark, MO 65721, (800) 321-2739

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Photographic Supply Bostick and Sullivan , PO Box 16639, Santa Fe, NM 87592, (505) 474-

0890 Photographer's Formulary , PO Box 950 Condon, MT 59826, (800) 922-

5255 Pyrotechnic Supply

Firefox , PO Box 5366, Pocatello, ID 83202, (208) 237-1976 Iowa Pyro Supply , 1000 130th Street, Stanwood, IA 52337, (563) 945-

6637 Skylighter , PO Box 480-W, Round Hill, VA 20142, (540) 554-4543 United Nuclear , PO Box 851 Sandia Park, NM. 87047 (505) 286-2831

Weaving Supply Handweavers Guild of America , (770) 495-7702 The Wool Co. , 990 US Highway 101, Bandon, OR 97411, (888) 456-2430 The Woolery , PO Box 468, Murfreesboro, NC 27855, (800) 441-9665 Paradise Fibers , 701 Parvin Road, Colfax WA 99111, (888) 320-7746

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Appendix E. Atomic Weights

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Appendix . InspirationsMost folks won't care about what's on this page,

but I think that the story's worth telling.The people here listed have helped shape the book.

Their bejeesical bits were compelling.

My mother and father were there from the startand imparted to me  my first memes.

Scoutmaster Bill Davidson helped me tie knotsand set fire to my kindling dreams.

Doug Boyd and Ed Guffee were buddies of minewith an int'rest in things that are old.

Together we dug up a lot of cool stufffor the LE Museum to hold.

Magnus McBride and the Mystery Schoolshowed me truth can be told with a lie.

Alain Nu and the Phoenix rose up from the ash;even memetic children must die.

Paul Mueller supported the Caveman I-deawhen most everyone thought I was mad.

Mary Prevo de-Og-lified some of the text,on account of the grammar was bad.

Five readers took time to set fire to the book;Brian Goeckerman and Aaron Skeen,

Cory Jaques and Bruce Miller played Athanor well,burning crap from the book with Ross Greene.

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The Cavemen of H-SC taught me a lotabout what even slackers can do.

I've loved every minute I've spent with you guys.I mean it—I wrote this for you.

And thanks to the "bots" (that's just Dunn-speak for cats),who insisted on helping me out.

Some of their typing was prob'ly left in.If you find it, just give me a shout.

And here's to my Sunshine, the light of my life,who put up with my preoccupation.

I love you for helping me finish this thing.Maybe now we can take a vacation.

Kevin Dunn is the Elliott Professor of Chemistry at Hampden-Sydney College, where he teaches the course which inspired this book. He holds a BS degree from the University of Chicago and a PhD from the University of Texas at Austin. He appears on The Learning Channel's Mysteries of Magic and is co-author of a dozen journal articles in theoretical chemistry. He lives in central Virginia with his wife and several cats.

This book was formatted from its DocBook SGML source using the open-source tools, Openjade, Jadetex, and LaTeX. The author would like to thank Norman Walsh, Sebastian Rahtz, Adam Di Carlo, and members of the DocBook-Apps mailing list.

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BibliographyAll quoted works or their English translations published prior to 1928 are assumed to be in the public domain. Permission has been granted by the respective publishers to reproduce those works protected by copyright. Reference [3] reprinted with the permission of MIT Press. Reference [6] courtesy of Hagley Museum and Library. Reference [23] reprinted by permission of the publishers and the Trustees of the Loeb Classical Library. The Loeb Classical Library is a registered trademark of the President and Fellows of Harvard College. Reference [25] reprinted with permission of the American Philosophical Society. References [53] and [65] reprinted with the permission of Cambridge University Press.

Brief snatches from literature, movies, television, and popular songs have been used to comment on the nature of memes under the fair use provision of the Copyright Act. Full bibliographic citations would be unwieldy; sources for these snatches have been listed in the index as a kind of key for those who recognize them.

Primary Texts

[1] Georgius Agricola, H. Hoover, L. Hoover, De Re Metallica, 1556, 1912, 1950, Dover Publications, Order from Amazon.com.

[2] Aristotle, The Works of Aristotle, Edited by W. D. Ross, 1908, Oxford University Press.

[3] Vannoccio Biringuccio, Cyril Smith, Martha Gnudi, The Pirotechnia, 1942, The M.I.T. Press, Order from Amazon.com.

[4] Robert Bradbury, An Inductive Chemistry, 1912, D. Appleton & Co..

[5] E. Caley, The Stockholm Papyrus, J. Chem. Educ., 3, 1927, 992-999.

[6] W. H. Carothers, Letter to Hamilton Bradshaw, November 9, 1927, Hagley Museum and Library, 1896.

[7] W. H. Carothers, J. A. Arvin, Studies of Polymerization and Ring Formation II Poly-esters, J. Am. Chem. Soc., 1929, American Chemical Society, 51, 2560.

[8] M. Chomell, The Dictionaire Oeconomique, 1758, L. Flinn.

[9] William Crookes, The Wheat Problem, 1899, G. P. Putnam's Sons.

Page 211of 234 Pages

[10] Jean D'Espagnet, 1623, 1650, 1987, Hermetic Arcanum, Holmes Publishing Group, Order from Amazon.com, Internet Sacred Text Archive.

[11] John Draper, Textbook on Chemistry, 1846, Harper & Brothers.

[12] Claudius Galenus, Opera Omnia, 2001, Georg Olms Verlag, Order from Amazon.com, Order from Amazon.com.

[13] L. W. King, Hammurabi, King of Babylonia, 1898, Internet Sacred Text Archive, Luzac & Co..

[14] Holy Bible (KJV), Internet Sacred Text Archive.

[15] Homer, Samuel Butler, The Iliad of Homer, 1898, Longmans, Green, and Co..

[16] Lao-tzu, Tao Te Ching, 1891, 1962, Internet Sacred Text Archive, Dover.

[17] Antoine-Laurent Lavoisier, Robert Kerr, Elements of Chemistry, 1790, 1965, Dover, Order from Amazon.com.

[18] Henry Leicester, Herbert Klickstein, Source Book in Chemistry, 1952, Harvard University Press.

[19] Lucretius, William Ellery Leonard, On the Nature of Things, 1916, E. P. Dutton, Internet Sacred Text Archive.

[20] Marie McLaughlin, Myths and Legends of the Sioux, 1916, Bismarck Tribune Company, Internet Sacred Text Archive.

[21] William Noyes, Textbook of Chemistry, 1913, Henry Holt & Co..

[22] Ralph Oesper, Christian Schonbein, J. Chem. Educ., 6, 1929, 432-440.

[23] Pliny, H. Rackham, Natural History, 1962, Harvard University Press, Order from Amazon.com.

[24] William Shakespeare, The Riverside Shakespeare, Houghton Mifflin, Order from Amazon.com.

[25] C. Smith, J. Hawthorne, Mappae Clavicula, 1974, American Philosophical Society, Order from Amazon.com.

[26] H. Stapleton, R. Azo, M. Husain, Chemistry in Iraq and Persia in the Tenth Century A.D., Mem. Asiat. Soc. Bengal, VIII, 1927, 317-418.

Page 212of 234 Pages

[27] Theophilus, J. Hawthorne, C. Smith, On Divers Arts, 1979, Dover Publications, Order from Amazon.com.

[28] Vitruvius, Morris Morgan, The Ten Books on Architecture, 1914, Harvard University Press, Order from Amazon.com.

[29] Rudolf Wagner, William Crookes, Wagner's Chemical Technology, 1872, 1988, Lindsay Publications.

Agencies and Organizations

[30] Agency for Toxic Substances and Disease Registry Website, 2002, U.S. Agency for Toxic Substances and Disease Registry.

[31] J. T. Baker Website, 2003, Mallinckrodt Baker, Inc..

[32] Bureau of Alcohol, Tobacco, and Firearms Website, 2002, U.S. Bureau of Alcohol, Tobacco, and Firearms.

[33] National Center for Chronic Disease Prevention and Health Promotion Website, 2003, National Center for Disease Prevention and Health Promotion.

[34] National Center for Injury Prevention and Control Website, 2003, National Center for Injury Prevention and Control.

[35] NFPA Website , 2003, National Fire Protection Association.

[36] Amy Spencer, Guy Colonna, 2002, Fire Protection Guide To Hazardous Materials, National Fire Protection Association, Order from Amazon.com.

[37] NPCA Website , 2003, National Paint &Coatings Association.

[38] OSHA Website , 2002, U.S. Occupational Safety and Health Administration.

[39] US Census Bureau Website , 2002, U.S. Census Bureau.

Brewing

[40] H. Bravery, Home Brewing Without Failures, 1965, Gramercy Publishing.

[41] Vincent Gingery, The Secrets of Building an Alcohol Producing Still, 1994, David Gingery, Order from Amazon.com.

[42] Lee Janson, Brew Chem 101, Order from Amazon.com, 1996, Storey Communications.

Page 213of 234 Pages

[43] Charles Papazian, The New Complete Joy of Homebrewing, Order from Amazon.com, 1991, Avon Books.

Chemical Demonstrations

[44] Bassam Shakhashiri, 1983, Chemical Demonstrations, V. 1, University of Wisconsin Press, Order from Amazon.com.

[45] Bassam Shakhashiri, 1985, Chemical Demonstrations, V. 2, University of Wisconsin Press, Order from Amazon.com.

Electricity

[46] Paul Doherty, The Aluminum/Air Battery, 2000, The Exploratorium.

[47] Alfred Morgan, The Boy Electrician, 1940, 1995, Lindsay Publications, Order from Amazon.com.

Fireworks

[48] Thomas Perigrin, Practical Introductory Pyrotechnics, 1996, Falcon Fireworks.

Glass

[49] Graham Stone, Firing Schedules for Glass, 1996, Graham Stone, Order from Amazon.com.

[50] Keith Cummings, Techniques of Kiln-Formed Glass, 1997, University of Pennsylvania, Order from Amazon.com.

[51] Jim Kervin, Dan Fenton, Pate de Verre and Kiln Casting of Glass, 1997, GlassWear Studios, Order from Amazon.com.

Industrial History

[52] Fred Aftalion, A History of the International Chemical Industry, 1991, University of Pennsylvania Press, Order from Amazon.com.

[53] Ahmad al-Hassan, Donald Hill, Islamic Technology, 1986, Cambridge University Press, Order from Amazon.com.

[54] Chemical and Engineering News, American Chemical Society, 0009-2347.

[55] Cathy Cobb, Harold Goldwhite, Creations of Fire, 1995, Plenum Press, Order from Amazon.com.

Page 214of 234 Pages

[56] Alfred Cowles, The True Story of Aluminum, 1958, Henry Regnery Company.

[57] Edited by Herman Daly, Kenneth Townsend, Valuing the Earth, 1993, Massachusetts Institute of Technology.

[58] F. W. Gibbs, The History of the Manufacture of Soap, Annals of Science, 1939.

[59] Simon Garfield, Mauve, 2002, W. W. Norton, Order from Amazon.com.

[60] L. F. Haber, The Chemical Industry During the Nineteenth Century, 1958, Clarendon Press.

[61] L. F. Haber, The Chemical Industry: 1900-1930, 1971, Clarendon Press.

[62] Matthew Hermes, Enough for One Lifetime, 1996, American Chemical Society, Order from Amazon.com.

[63] Joseph Lambert, Traces of the Past, 1997, Perseus Books, Order from Amazon.com.

[64] Charles Mann, Mark Plumer, The Aspirin Wars, 1991, Alfred Knopf, Order from Amazon.com.

[65] Joseph Needham, Science and Civilization in China: Chemistry and Chemical Technology: Military Technology, 5.7, 1986, Cambridge University Press, Order from Amazon.com.

[66] Joseph Needham, Science and Civilization in China: Chemistry and Chemical Technology: Spagyrical Discovery and Inventions, 5.4, 1980, Cambridge University Press, Order from Amazon.com.

[67] Joseph Needham, Science and Civilization in China: Chemistry and Chemical Technology: Paper and Printing, 5.1, 1985, Cambridge University Press, Order from Amazon.com.

[68] J. R. Partington, The Alkali Industry, 1925, Bailliere, Tindall, and Cox.

[69] James Partington, A History of Greek Fire and Gunpowder, 1960, 1999, The Johns Hopkins University Press, Order from Amazon.com.

[70] Naomi Rosenblum, A World History of Photography, 1997, Abbeville Press, Order from Amazon.com.

[71] Hugh Salzberg, From Caveman to Chemist, 1991, American Chemical Society.

[72] John Stillman, The Story of Early Chemistry, 1924, D. Appleton & Co..

Page 215of 234 Pages

[73] Graham Taylor, Patricia Sudnik, Du Pont and the International Chemical Industry, 1984, Twayne Publishers, Order from Amazon.com.

[74] Anthony Travis, The Rainbow Makers, 1993, Lehigh University Press, Order from Amazon.com.

[75] Kenneth Warren, Chemical Foundations: The Alkali Industry in Britain to 1926, 1980, Clarendon Press, Order from Amazon.com.

[76] Trevor Williams, A Biographical Dictionary of Scientists, 1969, Wiley-Interscience, Order from Amazon.com.

Memetics

[77] Richard Dawkins, The Selfish Gene, 1976, Oxford University Press, Order from Amazon.com.

[78] Susan Blackmore, The Meme Machine, 1999, Oxford University Press, Order from Amazon.com.

Paper

[79] Vance Studley, The Art and Craft of Handmade Paper, 1990, Dover Publications, Order from Amazon.com.

[80] Sophie Dawson, The Art and Craft of Papermaking, 1992, Running Press, Order from Amazon.com, Order from Amazon.com.

[81] Jules Heller, Paper-Making, 1997, Watson-Guptill, Order from Amazon.com.

Photography

[82] Edited by John Barnier, Coming into Focus, 2000, Chronicle Books, Order from Amazon.com.

[83] John Burk, Walter Henry, Paul Messier, Timothy Vitale, Albumen Photographs: History, Science, and Preservation, 2000, Stanford University.

[84] Randall Webb, Martin Reed, Alternative Photographic Processes, 2000, Silver Pixel Press, Order from Amazon.com.

Rocks and Minerals

[85] David Barthelmy, Mineralogy Database, 2002, Webmineral.com.

Page 216of 234 Pages

[86] Alan Holden, Phylis Morrison, Crystals and Crystal Growing, 1982, MIT Press, Order from Amazon.com.

[87] Tim Rast, Mike Melbourne, Knappers Anonymous Website, 2003, Mike Melbourne.

Soap

[88] Ann Bramson, Soap: Making It, Enjoying It, 1975, Workman Publishing, Order from Amazon.com.

[89] Susan Cavitch, The Soapmaker's Companion, 1997, Storey Publishing, Order from Amazon.com.

[90] Catherine Failor, Making Transparent Soap, 2000, Storey Publishing, Order from Amazon.com.

Textiles

[91] Rachel Brown, The Weaving, Spinning, and Dyeing Book, 1978, Alfred A. Knopf, Order from Amazon.com.

[92] J. N. Liles, The Art and Craft of Natural Dyeing, 1990, University of Tennessee Press, Order from Amazon.com.

Page 217of 234 Pages

GlossaryAcid

A substance which donates H+ (proton) to other substances.

Aerobic Process

A process which occurs in the presence of oxygen.

Air

That which is hot and moist. Alchemical air symbolizes daring and skill. In this book Air speaks for skepticism and honesty.

Alkali

A substance which accepts H+ (proton) from other substances. An alkali reacts with water to produce OH- in solution.

Anaerobic Process

A process which occurs in the absence of oxygen.

Base

An alkali.

Calcination

A procedure by which the bejeezus is driven from a substance, often by intense heat.

Catalyst

An agent which speeds up a reaction. A catalyst is not consumed in the reaction it catalyzes and so does not appear as a reactant or product in the balanced chemical equation.

Combustion

A process by which a material combines with oxygen, usually releasing heat.

Page 218of 234 Pages

Compound

A pure substance which can be decomposed into two or more other pure substances.

Condensation

A reaction in which two molecules join together, spitting out a small molecule, usually water, in the process.

Conservation of Mass

A generalization of the observation that the mass of a sealed container does not change.

Dissolution

A procedure by which one substance is dissolved in another to form a solution.

Distillation

A procedure by which the bejeezus is collected and isolated.

Earth

That which is cold and dry. Alchemical earth symbolizes wealth and bounty. In this book Earth speaks for business and industry.

Electrolyte

A substance which dissociates into ions when dissolved. Because these ions are free to move about, an electrolyte solution conducts electricity.

Element

A pure substance which has never been decomposed into two or more other pure substances.

Equation

A relationship between two things which are equal. In a chemical equation, the two things must have the same number and kinds of elements. In a mathematical equation, the two things must have the same units.

Page 219of 234 Pages

Experiment

A procedure for making an observation which will be consistent with only one of two or more theories.

Fermentation

A procedure by which one substance is converted to another by decay, often by the action of microorganisms.

Fire

That which is hot and dry. Alchemical fire symbolizes mastery and will. In this book Fire speaks for basic research and experimentation.

Formula

A symbol, e.g. H2O, used to denote the relative amounts of elements in a pure substance.

Glossary

The section of the book you are reading right now.

Heterogeneous Matter

Matter whose composition varies from one place to another within a sample.

Homogeneous Matter

Matter whose composition does not vary from one place to another within a sample.

Inorganic Compound

A compound which does not contain carbon. There are exceptions to this definition; carbonates, bicarbonates, and carbon dioxide are generally considered as inorganic compounds.

Lixiviation

A process by which water is passed over a solid, taking with it any soluble materials.

Page 220of 234 Pages

Meme

A unit of culture which is replicated through imitation. Successful memes exhibit fidelity, fecundity, and longevity. Synonyms include I-dea, inspiration, and spirit.

Mercury

That which flees the fire, descending from heaven to earth. Alchemical mercury symbolizes the spirit and takes its name from the element mercury, which may be distilled without leaving any ash. In this book Mercury provides the motivation for the Sulfur and Salt.

Metathesis

A procedure by which two substances exchange anions or cations.

Organic Compound

A compound which contains carbon. There are exceptions to this definition; carbonates, bicarbonates, and carbon dioxide are generally considered as inorganic compounds.

Oxidation

A half-reaction in which electrons appear as products.

pH

A measure of the hydrogen ion concentration in a solution. Solutions with pH less than 7 are acidic, which those with pH higher than 7 are alkaline.

Precipitation

A procedure by which a solid falls out of a liquid solution.

Recrystallization

A procedure for purifying a substance by precipitating a solid from a liquid.

Stoichiometric Coefficient

The number in front of each substance in a balanced chemical equation.

Page 221of 234 Pages

Stoichiometric Question

A question involving the relative weights of reactants and products in a chemical reaction.

Substance

Homogeneous matter whose composition is fixed.

Reduction

A half-reaction in which electrons appear as reactants.

Salt

That which remains from the fire, neither rising to heaven nor descending to earth. Alchemical salt symbolizes the body and takes its name from those salts which are extracted from ashes. In this book Salt provides instructions for making or doing something.

Separation

A procedure by which one substance is separated from another. Examples of separations include recrystallization, distillation, and chromatography.

Solute

The less abundant component of a solution.

Solution

Homogeneous matter whose composition is variable.

Solvent

The more abundant component of a solution.

Sublimation

A procedure for purifying a substance by precipitating a solid from a gas.

Page 222of 234 Pages

Sulfur

That which is consumed in the fire, rising from earth to heaven. Alchemical sulfur symbolizes the mind and takes its name from the element sulfur, which burns without leaving any ash. In this book Sulfur provides the chemical concepts needed to understand the Salt.

Theory

A logically consistent set of principles which account for, explain, or render intelligible a set of observations.

Water

That which is cold and moist. Alchemical water symbolizes belief and foresight. In this book Water speaks for theory and prediction.

Page 223of 234 Pages

IndexAbsorber, Spot and Roebuck (Acid), Solvay (Ammonia), Badische (Fertilizers)Acid, Adelard (Alcohol), Spot and Roebuck (Acid), Glossary

acetic, Samson (Mead), Adelard (Alcohol), Spot and Roebuck (Acid), Bath (Soap)fatty, Bath (Soap)saturated, Bath (Soap)unsaturated, Bath (Soap)nitric, Badische (Fertilizers), Badische (Fertilizers)oleic, Bath (Soap)organic, Bayer (Pharmaceuticals)palmitic, Bath (Soap)phthalic, DuPont (Plastics)stearic, Bath (Soap)sulfuric, Spot and Roebuck (Acid), Spot and Roebuck (Acid), Leblanc (Soda)contact process, Badische (Fertilizers)sulfurous, Spot and Roebuck (Acid)terephthalic, DuPont (Plastics)

Acid anhydride, Bayer (Pharmaceuticals)Aerobic process, GlossaryAir, Lucifer (Charcoal)

alchemical, GlossaryAlbumin, Eastman (Photography)Alcohol, Adelard (Alcohol), Bayer (Pharmaceuticals)

ethyl, Samson (Mead)Alkali, Job (Alkali), Job (Alkali), Ts'ai Lun (Paper), Spot and Roebuck (Acid), Leblanc (Soda), GlossaryAlloy, Vulcan (Metals)Alumina, Athanor (Ceramics)Aluminum, Dow (Electrochemicals)Amalgam, Vulcan (Metals)American Cyanamid, Badische (Fertilizers)Amide, Bayer (Pharmaceuticals)Amine, Perkin (Aniline?Dyes), Bayer (Pharmaceuticals)Ammonia, Marie (Dyes), Ts'ai Lun (Paper), Spot and Roebuck (Acid), Bath (Soap), Eastman (Photography), Solvay (Ammonia), Badische (Fertilizers), Badische (Fertilizers)Ammonium bicarbonate, Solvay (Ammonia)Ammonium chloride, Solvay (Ammonia)Ammonium nitrate, Badische (Fertilizers)Ampere, Volta (Batteries)Amp貥, Andr鬊 Volta (Batteries)Anaerobic process, GlossaryAnalogy, tail-sniffing, Adelard (Alcohol), Bath (Soap)

Page 224of 234 Pages

Andrews, JulieMy Favorite Things, Adam (Metathesis?Reactions)

Anhydrite, Vitruvius (Lime)Aniline, Perkin (Aniline?Dyes)Anion, Adam (Metathesis?Reactions)Annealing, Theophilus (Glass)Anode, Volta (Batteries)Aragonite, Vitruvius (Lime)Arlington Company, DuPont (Plastics)Aromatic compounds, Perkin (Aniline?Dyes)Aspdin, Joseph, Vitruvius (Lime)Aspirin, Bayer (Pharmaceuticals)Atlas Powder Company, DuPont (Plastics)Atomic weight, al-Razi (Stoichiometry)Azeotrope, Adelard (Alcohol)Bacteria, Job (Alkali), Marie (Dyes), Adelard (Alcohol), Spot and Roebuck (Acid)Badische Anilin und Soda Fabrik, Badische (Fertilizers)Baekeland, Leo, DuPont (Plastics)Bakelite, DuPont (Plastics)Balance, Measuring and Mixing

centigram, Measuring and Mixinggram, Measuring and Mixing

Barilla, Leblanc (Soda)Barth, Bergrat, Spot and Roebuck (Acid)Base, Job (Alkali), Job (Alkali), Adelard (Alcohol), Spot and Roebuck (Acid), GlossaryBassanite, Vitruvius (Lime)Battery, Volta (Batteries)

alkaline, Volta (Batteries)lead-acid, Dow (Electrochemicals)Leclanch鬊 Volta (Batteries)parallel, Volta (Batteries)series, Volta (Batteries)

Bayer Company, Bayer (Pharmaceuticals)Beddoes, Thomas, Dow (Electrochemicals)Bejeezus, Athanor (Ceramics), Vulcan (Metals), Vitruvius (Lime), Adelard (Alcohol), Perkin (Aniline?Dyes), Solvay (Ammonia), GlossaryBenzene, Perkin (Aniline?Dyes)Bevan, Edward, DuPont (Plastics)Beverly Hillbillies, Prologue, Badische (Fertilizers)Binder, Marie (Dyes)Black ash, Leblanc (Soda)Borax, Theophilus (Glass)Bosch, Carl, Badische (Fertilizers)Bots, feline, InspirationsBoyd, Doug, InspirationsBrass, Vulcan (Metals)

Page 225of 234 Pages

Bronze, Vulcan (Metals)Brooke, Walter

The Graduate, DuPont (Plastics)Browning, Robert

Song, From Pippa Passes, Leblanc (Soda)Burner, Lucifer (Charcoal), Spot and Roebuck (Acid), Badische (Fertilizers)Calcination, Athanor (Ceramics), GlossaryCalcite, Vitruvius (Lime)Calcium cyanamide, Badische (Fertilizers)California Powder Company, DuPont (Plastics)Camera obscura, Eastman (Photography)Carbon dioxide, Spot and Roebuck (Acid), Solvay (Ammonia)Caro, Nikodem, Badische (Fertilizers)Carothers, Wallace, DuPont (Plastics)Carpenter, Richard and Karen

We've Only Just Begun, Spot and Roebuck (Acid)CAS, Hammurabi (Units)Castaway, EpilogueCastner, Hamilton, Dow (Electrochemicals)Catalyst, Spot and Roebuck (Acid), Badische (Fertilizers), GlossaryCathode, Volta (Batteries)Cation, Adam (Metathesis?Reactions)Celanese Corporation, DuPont (Plastics)Celluloid, DuPont (Plastics)Cellulose, Lucifer (Charcoal), Venus (Textiles), Ts'ai Lun (Paper)Cellulose acetate, DuPont (Plastics)Charcoal, Tzu-Chhun (Gunpowder), Spot and Roebuck (Acid), Perkin (Aniline?Dyes)Chlorine, Dow (Electrochemicals)Chromatography, Lucifer (Charcoal), Perkin (Aniline?Dyes)Clay

body, Athanor (Ceramics)mineral, Athanor (Ceramics)

Clothing, Venus (Textiles)Coke, Perkin (Aniline?Dyes)Collodion, Eastman (Photography), Badische (Fertilizers)Colorfast dye, Marie (Dyes)Combustion, Lucifer (Charcoal), GlossaryCompound, Lucifer (Charcoal), GlossaryCondensation, Venus (Textiles), Adelard (Alcohol), Bath (Soap), Bayer (Pharmaceuticals), DuPont (Plastics), GlossaryCondenser, Adelard (Alcohol)Conjugate acid, Spot and Roebuck (Acid)Conjugate base, Spot and Roebuck (Acid)Conservation of Mass, Adam (Metathesis?Reactions), al-Razi (Stoichiometry), GlossaryCopper, Vulcan (Metals)Couche, Ts'ai Lun (Paper)

Page 226of 234 Pages

Coulomb, Charles, Volta (Batteries)Cross, Charles, DuPont (Plastics)Crystal, Unktomi (Silicates), Job (Alkali)Current, electric, Volta (Batteries)Dacron, DuPont (Plastics)Daguerre, Louis Jacques, Eastman (Photography)Davidson, Bill, InspirationsDavy, Humphry, Dow (Electrochemicals), Bayer (Pharmaceuticals)Deckle, Ts'ai Lun (Paper)Devitrification, Theophilus (Glass)Dickens, Charles, Leblanc (Soda)Dissolution, GlossaryDistillation, Lucifer (Charcoal), Glossary

destructive, Perkin (Aniline?Dyes), Perkin (Aniline?Dyes)Dogs, tail-sniffing, Adelard (Alcohol)Dolomite, Vitruvius (Lime)Dow, Herbert, Dow (Electrochemicals)Drafting, Venus (Textiles)Dreyfus, Henri and Camille, DuPont (Plastics)Drop spindle, Venus (Textiles)Du Pont de Nemours Powder Company, DuPont (Plastics)Du Pont, Irenee, Tzu-Chhun (Gunpowder), DuPont (Plastics)Duhamel, Henri Louis, Leblanc (Soda)Duisberg, Carl, Bayer (Pharmaceuticals)Dunn, James R., InspirationsDunn, Peggy, solar properties of, InspirationsDye, Marie (Dyes)Dynamite, Badische (Fertilizers)Earth, Lucifer (Charcoal), Unktomi (Silicates)

alchemical, GlossaryEastman Kodak, Bayer (Pharmaceuticals)Eastman, George, Eastman (Photography), DuPont (Plastics)Edison, Thomas, DuPont (Plastics)Electrolysis, Dow (Electrochemicals)Electrolyte, Adam (Metathesis?Reactions), Volta (Batteries), Dow (Electrochemicals), GlossaryElectromotive Force (EMF), Volta (Batteries)Element, Lucifer (Charcoal), GlossaryEmerald Tablet, Prologue, Hammurabi (Units), Vitruvius (Lime), Adelard (Alcohol), Tzu-Chhun (Gunpowder), Spot and Roebuck (Acid), Perkin (Aniline?Dyes), Solvay (Ammonia)Equation, Lucifer (Charcoal), Adam (Metathesis?Reactions), GlossaryErasmus, Desiderius, Badische (Fertilizers)Ester, Bath (Soap), Bayer (Pharmaceuticals), DuPont (Plastics)Ethanol, Samson (Mead)Experiment, Glossary

Page 227of 234 Pages

Experimental yield, al-Razi (Stoichiometry)Faraday, Michael, Volta (Batteries), Perkin (Aniline?Dyes), Dow (Electrochemicals), Bayer (Pharmaceuticals), DuPont (Plastics)Farbenfabriken vormals Friedrich Bayer, Bayer (Pharmaceuticals)Fat, Bath (Soap)Fecundity, Hammurabi (Units)Fermentation, Glossary

aerobic, Samson (Mead)anaerobic, Samson (Mead)

Fermenter, Samson (Mead)Fidelity, Hammurabi (Units)Figment, Hammurabi (Units), Tzu-Chhun (Gunpowder), Bath (Soap), Eastman (Photography), Bayer (Pharmaceuticals), Badische (Fertilizers), DuPont (Plastics), EpilogueFire, Lucifer (Charcoal), Lucifer (Charcoal)

alchemical, GlossaryFlux, Vulcan (Metals), Vulcan (Metals), Theophilus (Glass)Ford Motor Company, DuPont (Plastics)Formula, Lucifer (Charcoal), Adam (Metathesis?Reactions), Glossary

empirical, Athanor (Ceramics), Venus (Textiles)molecular, Venus (Textiles)oxide, Athanor (Ceramics)structural, Perkin (Aniline?Dyes)

Formula weight, al-Razi (Stoichiometry)Frank, Adolph, Badische (Fertilizers)Franklin, Benjamin, Volta (Batteries)Frit, Theophilus (Glass)Fugitive dye, Marie (Dyes)Furnace, Lucifer (Charcoal), Vitruvius (Lime), Adelard (Alcohol), Leblanc (Soda), Solvay (Ammonia), Badische (Fertilizers)Galvani, Luigi, Volta (Batteries)General Motors, DuPont (Plastics)Glass, Leblanc (Soda)Glauber, Rudolf, Spot and Roebuck (Acid), Leblanc (Soda), Leblanc (Soda)Glossary, Lucifer (Charcoal), Unktomi (Silicates), Hammurabi (Units), Samson (Mead), Athanor (Ceramics), Venus (Textiles), Adam (Metathesis?Reactions), Job (Alkali), Vulcan (Metals), Vitruvius (Lime), Pliny (Redox?Reactions), Marie (Dyes), Theophilus (Glass), Ts'ai Lun (Paper), al-Razi (Stoichiometry), Adelard (Alcohol), Tzu-Chhun (Gunpowder), Spot and Roebuck (Acid), Bath (Soap), Leblanc (Soda), Volta (Batteries), Perkin (Aniline?Dyes), Eastman (Photography), Solvay (Ammonia), Dow (Electrochemicals), Bayer (Pharmaceuticals), Badische (Fertilizers), DuPont (Plastics), GlossaryGlycerol, Bath (Soap)Goeckerman, Brian, InspirationsGoffin, Gerry

Who Put the Bomp, Hammurabi (Units)

Page 228of 234 Pages

Gold, Vulcan (Metals)Greene, Donald, InspirationsGuffee, Eddie, InspirationsGypsum, Vitruvius (Lime), Vitruvius (Lime)Haber, Fritz, Badische (Fertilizers)Half-reaction, Pliny (Redox?Reactions)Halite, Leblanc (Soda)Hall, Charles Martin, Dow (Electrochemicals)Hammerstein, Oscar

My Favorite Things, Adam (Metathesis?Reactions)Harrison Brothers, DuPont (Plastics)Hatcher, Margaret, InspirationsHe/his, Lucifer (Charcoal), Adam (Metathesis?Reactions), Adelard (Alcohol), Adelard (Alcohol), Spot and Roebuck (Acid), Leblanc (Soda), Volta (Batteries), Eastman (Photography), Badische (Fertilizers)Heat, Athanor (Ceramics)Henry, Buck

The Graduate, DuPont (Plastics)Hercules Powder Company, DuPont (Plastics)Heroult, Paul, Dow (Electrochemicals)Herschel, John, Eastman (Photography)Heterogeneous matter, Lucifer (Charcoal), Samson (Mead), GlossaryHofmann, Wilhelm, Perkin (Aniline?Dyes)Homogeneous matter, Lucifer (Charcoal), Samson (Mead), GlossaryHyatt, John, DuPont (Plastics)Hydraulic cement, Vitruvius (Lime)Hydrogen, Badische (Fertilizers)Index, Lucifer (Charcoal), Unktomi (Silicates), Hammurabi (Units), Samson (Mead), Athanor (Ceramics), Venus (Textiles), Adam (Metathesis?Reactions), Job (Alkali), Vulcan (Metals), Vitruvius (Lime), Pliny (Redox?Reactions), Marie (Dyes), Theophilus (Glass), Ts'ai Lun (Paper), al-Razi (Stoichiometry), Adelard (Alcohol), Tzu-Chhun (Gunpowder), Spot and Roebuck (Acid), Bath (Soap), Leblanc (Soda), Volta (Batteries), Perkin (Aniline?Dyes), Eastman (Photography), Solvay (Ammonia), Dow (Electrochemicals), Bayer (Pharmaceuticals), Badische (Fertilizers), DuPont (Plastics)Indigo, Spot and Roebuck (Acid)Indigo carmine, Spot and Roebuck (Acid)Indigotin, Marie (Dyes)Inorganic compound, Adam (Metathesis?Reactions), GlossaryInvestment, Theophilus (Glass)Ionic substance

See Substance, ionicIonization, Adam (Metathesis?Reactions)Iron, Vulcan (Metals)Isn't It Grand, Boys?, Pliny (Redox?Reactions)Isomer, Perkin (Aniline?Dyes), DuPont (Plastics)Jaques, Cory, Inspirations

Page 229of 234 Pages

John, EltonCandle in the Wind, DuPont (Plastics)

Jordan, LouisAin't Nobody Here But Us Chickens, Epilogue

Kaye, Gorden'Allo 'Allo!, Prologue

Kellner, Karl, Dow (Electrochemicals)Kelp, Leblanc (Soda)Kolbe, Hermann, Bayer (Pharmaceuticals)Lavigne, Avril

Complicated, al-Razi (Stoichiometry)Lavoisier, Antoine, Lucifer (Charcoal), Tzu-Chhun (Gunpowder), Dow (Electrochemicals), DuPont (Plastics)LD50, Adam (Metathesis?Reactions)Lead, Vulcan (Metals)Leblanc, Nicolas, Leblanc (Soda)Lebon, Phillipe, Perkin (Aniline?Dyes)Led Zeppelin

Stairway to Heaven, Perkin (Aniline?Dyes)Leone, Sergio

The Good, the Bad, and the Ugly, Badische (Fertilizers)Lightfoot, Gordon

If You Could Read My Mind, PrologueLignin, Ts'ai Lun (Paper)Lime, Vitruvius (Lime), Theophilus (Glass), Ts'ai Lun (Paper), Spot and Roebuck (Acid), Bath (Soap), Leblanc (Soda)

quick, Vitruvius (Lime)slaked, Vitruvius (Lime), Ts'ai Lun (Paper), al-Razi (Stoichiometry)

Limestone, Vitruvius (Lime), Vitruvius (Lime)Limiting reagent, al-Razi (Stoichiometry), Solvay (Ammonia)Lincoln, Abraham

Gettysburg Address, Solvay (Ammonia)Lister, Joseph, Bayer (Pharmaceuticals)Liston, Robert, Bayer (Pharmaceuticals)Litharge, Theophilus (Glass)Lixiviation, Leblanc (Soda), GlossaryLixiviator, Job (Alkali), Leblanc (Soda)Logan, Horace, EpilogueLongevity, Hammurabi (Units)Lye, al-Razi (Stoichiometry), Bath (Soap)Maceration, Ts'ai Lun (Paper)Mann, Barry

Who Put the Bomp, Hammurabi (Units)Manure, Spot and Roebuck (Acid)Martin, Steve

The Jerk, Ts'ai Lun (Paper)

Page 230of 234 Pages

McBride, Jeff, InspirationsMcLean, Don

Vincent, DuPont (Plastics)Mead, Samson (Mead)Mellville, Herman

Moby Dick, Lucifer (Charcoal)Meme, Hammurabi (Units), GlossaryMercury, Lucifer (Charcoal), Vulcan (Metals), Adelard (Alcohol), Tzu-Chhun (Gunpowder), Dow (Electrochemicals)

alchemical, GlossaryMetathesis, GlossaryMiller, Bruce, InspirationsMineral, Unktomi (Silicates)Mixture, Samson (Mead)Mold, Ts'ai Lun (Paper)Mole, Adam (Metathesis?Reactions)Molecular model, Perkin (Aniline?Dyes)Molecular weight, al-Razi (Stoichiometry)Monomer, Venus (Textiles)Mordant, Marie (Dyes)Morse, Samuel, Volta (Batteries)Mortimer, John, EpilogueMSDS, Hammurabi (Units)Mueller, Paul, InspirationsMurdock, William, Perkin (Aniline?Dyes)Murrah Federal Building, Badische (Fertilizers)Napoleon, Leblanc (Soda)Napoleon III, Eastman (Photography), Dow (Electrochemicals)Napoleon XIV

They're Coming to Take Me Away, Badische (Fertilizers)Natron, Leblanc (Soda)Negative, photographic, Eastman (Photography)Neoprene, DuPont (Plastics)NFPA Diamond, al-Razi (Stoichiometry)Nieuwland, Julius, DuPont (Plastics)Nitrocellulose, Badische (Fertilizers), Badische (Fertilizers), DuPont (Plastics)Nitrogen, Lucifer (Charcoal), Badische (Fertilizers)Nitroglycerin, Badische (Fertilizers), Badische (Fertilizers)Nobel, Alfred, Badische (Fertilizers), DuPont (Plastics)Non-polar substance

See Substance, non-polarNose, chemist's virtual, Leblanc (Soda)Nu, Alain, InspirationsNylon, DuPont (Plastics)Oersted, Hans Christian, Volta (Batteries), Dow (Electrochemicals)Ohm, Georg, Volta (Batteries)

Page 231of 234 Pages

Oil, Bath (Soap)Ore, Vulcan (Metals)Organic compound, Adam (Metathesis?Reactions), Adelard (Alcohol), GlossaryOrleans, Duke of, Leblanc (Soda)Ostwald, Wilhelm, Badische (Fertilizers)Oxidation, Athanor (Ceramics), Pliny (Redox?Reactions), Tzu-Chhun (Gunpowder), GlossaryOxygen, Lucifer (Charcoal), Badische (Fertilizers)Palahniuk, Chuck

Fight Club, Bath (Soap)Paper, Ts'ai Lun (Paper), Leblanc (Soda)Paper Lace

The Night Chicago Died, EpiloguePapyrus, Ts'ai Lun (Paper)Parchment, Ts'ai Lun (Paper)Parkes, Alexander, DuPont (Plastics)Part, Hammurabi (Units), Vitruvius (Lime), Theophilus (Glass), al-Razi (Stoichiometry), Spot and Roebuck (Acid)Perkin, William, Perkin (Aniline?Dyes), Bayer (Pharmaceuticals)Pewter, Vulcan (Metals)pH, Job (Alkali), GlossaryPhenol, Bayer (Pharmaceuticals)Pigment, Marie (Dyes)Pink Floyd

Another Brick in the Wall, Vitruvius (Lime)Pitt, Brad

Fight Club, Bath (Soap), Bath (Soap)Plaster, Vitruvius (Lime), Vitruvius (Lime)Ply, Venus (Textiles)Point

annealing, Theophilus (Glass)softening, Theophilus (Glass)strain, Theophilus (Glass)working, Theophilus (Glass)

Poise, Theophilus (Glass)Polar substance

See Substance, polarPolymer, Venus (Textiles)Post, Ts'ai Lun (Paper)Potash, Job (Alkali), Vulcan (Metals), Theophilus (Glass), Ts'ai Lun (Paper), Ts'ai Lun (Paper), Spot and Roebuck (Acid), Bath (Soap), Leblanc (Soda)

caustic, Bath (Soap)Potassa, Theophilus (Glass)Potassium, Dow (Electrochemicals)Precipitation, Adam (Metathesis?Reactions), GlossaryPrevo, Mary, Inspirations

Page 232of 234 Pages

Prine, JohnYou Got Gold, Adam (Metathesis?Reactions)

Problematic unit, Hammurabi (Units), al-Razi (Stoichiometry)Pulp, Ts'ai Lun (Paper)

chemical, Ts'ai Lun (Paper)mechanical, Ts'ai Lun (Paper)

Rayon, DuPont (Plastics)Reaction

acid-base, Adam (Metathesis?Reactions), Adelard (Alcohol)hydrolysis, Ts'ai Lun (Paper), Bath (Soap)metathesis, Adam (Metathesis?Reactions)

Reactorabsorber, Spot and Roebuck (Acid)burner, Lucifer (Charcoal)fermenter, Samson (Mead)furnace, Lucifer (Charcoal)lixiviator, Job (Alkali), Leblanc (Soda)slaker, Vitruvius (Lime)still, Adelard (Alcohol)

Recrystallization, Lucifer (Charcoal), Job (Alkali), GlossaryReduction, Athanor (Ceramics), Vulcan (Metals), Pliny (Redox?Reactions), Tzu-Chhun (Gunpowder), GlossaryRefractory, Athanor (Ceramics)Rock, Unktomi (Silicates)

igneous, Unktomi (Silicates)metamorphic, Unktomi (Silicates)sedimentary, Unktomi (Silicates), Vitruvius (Lime)

Roebuck, John, Spot and Roebuck (Acid)Rogers, Richard

My Favorite Things, Adam (Metathesis?Reactions)Runge, Friedlieb, Perkin (Aniline?Dyes)Russell, Bobby

The Night the Lights Went Out in Georgia, EpilogueSalt, Adam (Metathesis?Reactions), Adelard (Alcohol), Tzu-Chhun (Gunpowder), Spot and Roebuck (Acid)

alchemical, GlossaryGlauber's, Spot and Roebuck (Acid), Leblanc (Soda)sea, Leblanc (Soda)

Saltpeter, Tzu-Chhun (Gunpowder), Spot and Roebuck (Acid), Badische (Fertilizers)Chile, Badische (Fertilizers)

Saltwort, Leblanc (Soda)Sanford and Son, EpilogueSaponification, Bath (Soap)Saxony blue, Spot and Roebuck (Acid)Schonbein, Christian, Badische (Fertilizers), DuPont (Plastics)Separation, Glossary

Page 233of 234 Pages

She/her, Lucifer (Charcoal), Hammurabi (Units), Hammurabi (Units), Adam (Metathesis?Reactions), Adam (Metathesis?Reactions), Pliny (Redox?Reactions), Adelard (Alcohol), Tzu-Chhun (Gunpowder), Tzu-Chhun (Gunpowder), Perkin (Aniline?Dyes), Badische (Fertilizers)Silica, Unktomi (Silicates), Athanor (Ceramics)Silver, Vulcan (Metals)Simon, Paul

Bridge Over Troubled Waters, Volta (Batteries)Kodachrome, Eastman (Photography)Sound of Silence, DuPont (Plastics)

Sizing, Eastman (Photography)Skeen, Aaron, InspirationsSlaker, Vitruvius (Lime), Solvay (Ammonia)Smelting, Vulcan (Metals)SmithKline Beecham, Bayer (Pharmaceuticals)Soap, Leblanc (Soda)Sobrero, Ascanio, Badische (Fertilizers)Soda, Theophilus (Glass), Ts'ai Lun (Paper), Spot and Roebuck (Acid), Solvay (Ammonia)

ash, Job (Alkali), Vulcan (Metals), Theophilus (Glass), Ts'ai Lun (Paper), al-Razi (Stoichiometry), Bath (Soap), Leblanc (Soda)caustic, Ts'ai Lun (Paper), al-Razi (Stoichiometry), Spot and Roebuck (Acid), Bath (Soap), Leblanc (Soda), Dow (Electrochemicals)washing, al-Razi (Stoichiometry), Volta (Batteries)

Sodium, Dow (Electrochemicals)Sodium bicarbonate, Solvay (Ammonia)Sodium carbonate, Solvay (Ammonia)Sodium chloride, Solvay (Ammonia)Sodium hyposulfite, Eastman (Photography)Sodium thiosulfate, Eastman (Photography)Solder, Vulcan (Metals)Solubility, Adam (Metathesis?Reactions)Solute, Samson (Mead), Adam (Metathesis?Reactions), Theophilus (Glass), Adelard (Alcohol), GlossarySolution, Lucifer (Charcoal), Samson (Mead), GlossarySolvay, Ernest, Solvay (Ammonia)Solvent, Samson (Mead), Adam (Metathesis?Reactions), Theophilus (Glass), Adelard (Alcohol), GlossarySpider, Itsy-Bitsy, Unktomi (Silicates)Spinning, Venus (Textiles)Spirit lamp, Adelard (Alcohol), Adelard (Alcohol), Spot and Roebuck (Acid), Leblanc (Soda), Solvay (Ammonia)Sterling Products, Bayer (Pharmaceuticals)Steve Miller Band

Fly Like an Eagle, EpilogueStill, Adelard (Alcohol), Solvay (Ammonia)

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Stoichiometric coefficient, Adam (Metathesis?Reactions), Pliny (Redox?Reactions), al-Razi (Stoichiometry), GlossaryStoichiometric question, al-Razi (Stoichiometry), Tzu-Chhun (Gunpowder), Tzu-Chhun (Gunpowder), Bath (Soap), Leblanc (Soda), GlossaryString, Venus (Textiles)Sublimation, GlossarySubstance, Lucifer (Charcoal), Glossary

ionic, Adelard (Alcohol), Bath (Soap)non-polar, Adelard (Alcohol), Bath (Soap)polar, Adelard (Alcohol), Bath (Soap)

Sucrose, Tzu-Chhun (Gunpowder)Sugar, Samson (Mead)Sulfur, Tzu-Chhun (Gunpowder), Tzu-Chhun (Gunpowder), Spot and Roebuck (Acid)

alchemical, GlossarySulfur dioxide, Spot and Roebuck (Acid)Sulfur trioxide, Spot and Roebuck (Acid)Swan, Joseph, DuPont (Plastics)Talbot, Fox, Eastman (Photography)Temperature, Athanor (Ceramics)Theoretical yield, al-Razi (Stoichiometry), Leblanc (Soda), Solvay (Ammonia)Theory, Samson (Mead), GlossaryThomas, Larry

Seinfeld's Soup Nazi, Bath (Soap)Tin, Vulcan (Metals)Toluene, Perkin (Aniline?Dyes)Toluidine, Perkin (Aniline?Dyes)Tongue, chemist's virtual, Job (Alkali), Spot and Roebuck (Acid)Toxicity

acute, Adam (Metathesis?Reactions)chronic, Pliny (Redox?Reactions)

Tri-glyceride, Bath (Soap)Trinitrotoluene, Badische (Fertilizers)Trona, Leblanc (Soda)Twine, Venus (Textiles)Twist, s and z, Venus (Textiles)Uhls, Jim

Fight Club, Bath (Soap)Union Carbide, Badische (Fertilizers)Unit, Hammurabi (Units)Unit factor, Hammurabi (Units)Unit Factor Analysis (UFA), Hammurabi (Units)Unverdorben, Otto, Perkin (Aniline?Dyes)Urine, Spot and Roebuck (Acid)Vellum, Ts'ai Lun (Paper)Venus de Lespugue, Venus (Textiles)Vincenzoni, Luciano, Badische (Fertilizers)

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Vinegar, Samson (Mead), Adelard (Alcohol)Viscosity, Theophilus (Glass)Vitrification, Athanor (Ceramics)Vitriol, Spot and Roebuck (Acid)Volatile, Adelard (Alcohol)Volt, Volta (Batteries)Volta, Allesandro, Volta (Batteries)Voltaic cell, Volta (Batteries)von Liebig, Justus, Bayer (Pharmaceuticals)Ward, Joshua Spot, Spot and Roebuck (Acid)Water, Lucifer (Charcoal), Badische (Fertilizers)

alchemical, GlossaryWeighing by difference, Vitruvius (Lime), Theophilus (Glass), Tzu-Chhun (Gunpowder), Measuring and MixingWeiss, William, Bayer (Pharmaceuticals)Weston, Edward, DuPont (Plastics)Whorl, Venus (Textiles)Willingham, Calder

The Graduate, DuPont (Plastics)Wilson, Richard

One Foot in the Grave, Hammurabi (Units)Wood, Lucifer (Charcoal)Yeasts, Samson (Mead)Zevon, Warren

The Indifference of Heaven, Epilogue

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